Laboratory  Directions 


IN 


EDITED   BY 

WILLIAM    C.    BRAY,    PH.  D. 
ASSISTANT  PROFESSOR  OF  CHEMISTRY 

IN  THE 

UNIVERSITY    OF    CALIFORNIA 
1915 


Copyright  applied  for  by  William  C.  Bray. 


Lederer,   Street   &   Zeus   Co.,   Printers 


c; 


CHEMISTRY  lA.     UNIVERSITY,  0? '•„  EA-LjIOIOfiA' ::.-' 

LABORATORY  DIRECTIONS,  1915 

Provide  yourself  with  a  notebook  about  8  inches  by  5  inches  in  size  and 
not  loose-leaved. 

Notes  about  recording  the  results  of  experiments  are  given  in  the  first 
three  assignments  and  these  notes  apply  to  all  later  assignments. 

Bring  your  lecture  notes  and  textbook  of  inorganic  chemistry  into  the 
laboratory  and  use  them  for  reference. 

Satisfactory  progress  can  be  made  in  this  course  only  when  the  student 
accepts  the  responsibility  of  mastering,  before  the  next  period,  the  material 
presented  in  each  lecture  and  laboratory  period.  Consult  your  instructor  in  the 
laboratory  or  at  the  quiz  section  about  your  difficulties.  Frequent  brief  tests 
will  be  given,  usually  at  the  beginning  of  the  period. 

At  the  first  meeting  of  your  laboratory  section  go  to  the  room  to  which 
you  have  been  asigned  (consult  posted  lists  or  inquire  at  office  if  necessary), 
apply  to  your  instructor  for  your  key,  check  your  apparatus  by  means  of  the 
list  in  your  locker,  sign  your  name  (surname  first)  in  the  proper  place,  and 
return  this  list  to  your  instructor.  Begin  the  first  assignment. 

If,  while  checking  your  apparatus  during  the  first  laboratory  period,  any 
article  is  found  to  be  damaged,  exchange  it  at  the  office,  or  sign  a  blue  "return 
slip"  in  order  that  you  may  receive  proper  credit  if  the  damaged  article  is 
broken  during  the  term. 

Fill  your  wash-bottle  with  distilled  water.  Sterilize  the  mouth-piece  by 
placing  it  in  boiling  water  for  a  minute  or  two.  Always  use  the  distilled  water 
in  your  wash-bottle  for  the  final  rinsing  of  apparatus. 

Rinse  out  your  five  reagent  bottles,  partially  fill  them  with  the  6  normal 
laboratory  solutions.  Note  that,  while  pouring  the  reagent  from  one  of  these 
bottles,  the  stopper  of  the  bottle  should  be  held  with  the  same  hand  that  holds 
the  test  tube  or  beaker. 

In  general,  avoid  contamination  of  any  of  the  reagents  in  the  laboratory 
by  keeping  each  stopper  clean  and  returning  it  at  once  to  the  proper  bottle. 

Articles  to  replace  broken  apparatus,  and  additional  articles  for  regular 
use,  are  obtained  at  the  office  in  the  corridor  (on  the  same  floor  as  your  room) 
by  filling  out  a  white  order  slip  and  signing  your  name  and  desk  number. 

A  yellow  "temporary  order  slip"  is  used  instead  of  a  white  order  slip  when 
borrowing  articles  specified  in  the  last  paragraph  of  the  list  of  apparatus 
given  below.  These  articles  must  be  returned  during  the  same  laboratory  period 
before  the  office  is  closed. 

Sign  a  blue  "return  slip"  whenever  any  article  is  returned  to  the  office. 
The  article  will  not  be  accepted  unless  it  is  clean  and  in  good  condition. 

List  of  Apparatus. 

I.  Regular  equipment  of  each  locker.  At  the  end  of  the  term  the  locker 
must  contain  the  same  amount  of  apparatus,  no  more  and  no  less ;  the  locker 
must  be  clean;  the  apparatus  must  be  clean  and  dry,  and  in  good  condition; 
glass  stoppers  must  fit,  and  be  protected  by  the  insertion  of  a  piece  of  paper. 

1  Key. 

.5     Beakers,    100  cc.,    150  cc.,   250  cc.,   400  cc.,   600  cc. 
5     Reagent    Bottles. 

2  Sample   Bottles,    50  cc. 

i     Graduated   Cylinder,   50  cc. 

4     Flasks,    500   cc.,    250   cc.,    and   two    125    cc. 

i     Wash-bottle,    equipped   with   glass    tubing   and    rubber   stopper. 

314^92 


2     Blue  glasses. 

2     Glass  Rods,   12  cm.  and   18  cm. 
30     cm.   Glass  Tubing. 
12     Test-tubes. 

i     Watch  Glass. 

i     Casserole. 

1  Crucible,  with  cover. 

2  Evaporating  Dishes,  6  cm.  and  9  cm. 
15     cm.  Rubber  Tubing,  3  mm. 

i     Bunsen   Burner,  with  rubber  tubing. 

i     Wire  Gauze. 

i     Triangle. 

i     Test-tube  Brush. 

i     Test-tube  Holder. 

i     Test-tube  Rack. 

i     Package  Filter  Paper. 

1  Rule. 

2  Towels. 

II.     Additional  articles  may  be  obtained  at  the  office: 

(a)  By  signing  the   regular  white  order  slips. 
Special  apparatus  for  Assignments  III  and  V,  first  term. 
Matches. 

Glass  Flasks,  50  cc. 
Corks. 

Rubber  Stoppers. 

Platinum  Wire  and   Platinum  Foil.   Partial  credit  will  be  given  at  the 
end  of  each  term  for  Platinum  Wire  and  Foil  returned  to  the  office. 

(b)  By  signing  yellow  temporary  order  slips. 
Burettes,  with  clamps  and  pinchcocks. 
Graduated  Cylinders,  10  cc.  and  250  cc. 

Files. 
Paraffin. 

ASSIGNMENT    I. 
MEASUREMENT  OF  VOLUMES. 

Notes:  Enter  the  date  and  the  title  of  the  exercise  at  the  top  of  a  page 
of  the  notebook.  Use  one  side  only  of  each  page  for  recording  experimental 
results  and  answering  questions.  Make  calculation^  neatly  at  the  bottom  or 
side  of  the  page,  or  on  the  opposite  page.  Show  your  notebook  to  your 
instructor  each  day  until  he  is  satisfied  with  your  method  of  making  entries. 

Each  day,  before  leaving  the  laboratory;  finish  the  experiments  assigned 
for  the  day's  work,  write  out  the  answers  to  as  many  of  the  questions  as  pos- 
sible, and  clean  the  top  of  your  desk. 

Before  the  next  laboratory  period  finish  the  questions  and  problems, 
review  your  work  and  be  sure  that  you  understand  it  perfectly.  Read  over  the 
laboratory  notes  on  the  next  laboratory  exercise. 

Experiment.  Make  several  measurements  of  the  inside  diameter  (in 
centimeters)  of  your  medium-sized  beaker.  (The  instructor  will  suggest  a 
method.)  Record  each  measurement  in  your  notebook  as  soon  as  it  is 
made.  The  beakers  are  not  perfectly  cylindrical,  nor  is  the  method  very 
accurate ;  the  measurements,  therefore,  will  usually  show  variations  of  a  few 
tenths  of  a  centimeter.  An  average  of  at  least  four  measurements  will  give  the 
mean  diameter  to  about  o.i  cm.  Questions:  What  fraction  of  the  diameter 
is  this  error?  What  percent  of  the  diameter  is  this  error? 

Experiment.     Fill  the  beaker  with  water  to  just  belowr  the  flare.     Measure 


the  depth  of  the  water  by  dipping  the  rule  into  the  beaker.  Repeat  the  meas- 
urement. Mark  the  level  of  the  water  by  pasting  a  piece  of  gummed  paper  on 
the  outside  of  the  beaker.  Do  not  pour  the  water  out  of  the  beaker. 

Calculate  the  area  of  the  cross  section  of  the  beaker  from  your  average 
value  of  the  diameter  (assume  pi  =  3.14).  From  this  result  and  the  depth  of 
the  water,  calculate  the  volume  of  the  water  in  the  beaker. 

Make  further  calculations  to  determine  (i)  what  the  change  in  the  area 
of  the  cross  section  of  the  beaker  would  have  been  (in  sq.  cm)  if  the  diameter 
had  been  o.i  cm.  greater  than  the  value  chosen;  and  (2)  what  the  change  in 
the  calculated  volume  (in  cc.)  would  have  been  if  the  depth  of  the  water  had 
also  been  o.i  cm.  greater.  Questions:  What  percentage  error  in  the  area  of 
the  cross  section  results  from  an  error  of  o.i  cm.  in  the  diameter?  What 
percentage  error  in  the  volume  results  from  errors  of  o.i  cm.  in  the  diameter 
and  o.i  in  the  depth?  How  many  decimal  places  should  be  used  in  expressing 
the  final  result  for  the  volume? 

Experiment.  Measure  the  volume  of  water  in  the  beaker  used  in  the 
previous  experiment  by  means  of  a  measuring  cylinder  graduated  in  cc. 

Compare  this  measured  volume  with  the  volume  calculated.  If  the  diff- 
erence is  greater  than  the  estimated  error,  repeat  the  calculations  and  study 
the  following  discussion  to  determine  the  reason.  Repeat  the  measurements  if 
necessary. 

Note.  If  you  discover  an  error  in  your  previous  work,  do  not  erase  your 
first  entry,  nor  tear  out  the  page.  Mark  the  rejected  result  and  insert  the 
correction  near  the  original  entry. 

Note  that  the  central  surface  of  the  water  in  a  large  beaker  is  level.  Show 
in  a  drawing  the  appearance  of  a  vertical  section  of  the  surface  of  the  water 
(i)  where  it  touches  the  glass  in  the  large  beaker,  and  (2)  in  a  small  glass 
tube.  How  great  is  the  difference  (in  cc.)  between  the  readings  at  the  top  and 
the  bottom  of  the  meniscus  in  your  measuring  cylinder?  Volume  readings  for 
water  in  glass  vessels  are  made  at  the  bottom  of  the  meniscus. 

Compare  your  notebook  with  the  following  sample  page.  Note  whether 
you  have  arranged  your  notes  neatly  and  systematically. 

Measurement  of  Volumes.       Aug 1915 

Used  3rd   largest  beaker. 

Diameter,     6.3  cm.  Depth    of    water,    7.2  cm. 

6.4     "  7-3    " 

6.3     "  7-2    ' 

6.2 

Average,  7.2 

Average,       6.3 

o.i   is  0.1/6.3  =  0.016  of  the  diameter.  This  is % 

of    the    diameter. 

Area  of  cross  section    (dia.  6.3   cm.)     =  31.2  sq  cm. 
Volume  calculated   (depth  7.2  cm.)         =  225    cc. 
Area  calculated   (dia.  6.4  cm.)  =  32.2   sq.    cm. 

Difference,       =     i.o  sq.   cm.     i.  e %of  the  area. 

Volume    calculated    (dia.    6.4    cm,  =  cc. 

depth  7.3  cm.) 

Difference,       =  cc.,    i.  e %    of   the  volume. 

Volume  measured  with  graduate  =  219  cc. 

Difference,       =       6    cc.    less. 
This  result  .is   (or  is  not?)   within  the  experimental  error. 

Experiments.      Measure    by   means   of   a   measuring   cylinder   the   volumes 

5 
& 

(bfMfw    -  „ 

^  _    o 


of  a  second  beaker,  a  test-tube,  and  a  flask.  As  an  exercise  in  judging  vol- 
umes, estimate  the  volume  of  another  beaker  and  then  measure  it.  Record  in 
a  table  the  volumes  of  your  flasks  and  beakers.  What  is  the  volume  in  liters 
of  your  largest  flask? 

Experiments.  Pour  10  cc.  of  water  into  a  small  test  tube.  Mark  the  height 
of  the  water  by  pasting  a  small  label  on  the  side  of  the  tube.  Graduate  this 
test  tube  in  the  same  way  for  5  cc.  and  2  cc.,  and  set  it  aside  for  use  in  future 
experiments.  Graduate  your  smallest  beaker  for  volumes  of  50  cc.  and  100 
cc.  in  the  same  way. 

The  paper  labels  on  the  reagent  bottles  in  the  laboratory  are  covered  with 
paraffin.  Why?  Apply  to  the  instructor  to  have  the  labels  on  your  graduated 
test  tube  and  beaker  paraffined  in  a  similar  manner. 

State  how  you  would  use  your  graduated  cylinder  to  measure  out  75 
grams  of  water. 

Experiment,  to  give  practice  in  weighing.  Weigh  to  O.I  g.  the  small  beaker 
which  you  have  graduated.  (See  Notes  on  Weighing  at  the  end  of  Assign- 
ment). Fill  the  beaker  to  the  50  cc.  mark,  and  weigh.  What  is  the  weight  of 
water  in  the  beaker?  Calculate  its  volume  from  this  result. 

Problems.  (Always  answer  the  problems  before  the  next  laboratory 
period.) 

1.  How  much  water  would  you  weigh  out  in  order  to  have  a  volume  of 
10  cc.  ? 

2.  How  much  mercury  should  be  weighed  out  to  give  a  volume  of   10 
cc.?     The  density  of  mercury  is  13.6. 

3.  What  is  the  weight  of  a  liter  of  air?     The  density  of  air  is  0.0012  at 
room  temperature  and  atmospheric  pressure. 

4.  Mention    three    methods    of    determining    the    volume    of    water    in    a 
cylindrical   beaker. 

Bring  a  table  of  atomic  weights  (e.  g.  Cady,  page  96)  to  the  laboratory 
for  use  in  Assignments  II-V. 

Note  on  Weighing.  Follow  carefully  the  directions  given  by  the  instruc- 
tor. Report  to  him  at  once  if  the  balance  is  out  of  order  or  if  a  weight  is 
missing.  Enter  the  results  of  each  weighing  in  your  notebook  while  the 
weights  are  on  the  scale-pan  and  check  this  result  while  you  are  returning  the 
weights  to  their  places  in  the  box.  Summarize  in  a  table  the  results  of  the 
different  weighings  in  each  experiment. 

ASSIGNMENT     II. 

THE    SYNTHESIS    OF    COPPER    SULFIDE.      WEIGHT    RELATIONS    IN 
CHEMICAL  REACTIONS. 

(Enter    the    date    and    title    of    this    experiment    on    a    new    page.) 

The  purpose  of  this  experiment  is  to  determine  the  weights  of  copper  and 

sulfur  which  unite  to   form  a   sulfide  of  copper,   and  to  calculate  its   formula 

with  the  aid  of  the  atomic  weights  of  copper  and  sulfur. 

Experiment.  Support  a  clean  porcelain  crucible,  with  a  cover,  on  a  tri- 
angle and  heat  with  the  colorless  flame  of  a  bunsen  burner  to  low  redness.  Let 
the  crucible  cool  about  15  minutes,  and  weigh  it,  with  the  cover,  to  10  milli- 
grams. The  weights  and  the  balances  are  such  that  weighings  cannot  be 
made  closer  than  10  mg. 

While  the  crucible  is  cooling  obtain  a  clean  piece  of  copper  wire,  weigh- 
ing about  i  gram,  from  the  shelf  and  weigh  it  to  10  mg.  Question:  If  the 
weight  of  the  copper  is  in  error  by  10  mg.,  what  is  the  percentage  error? 

Place  the  copper  in  the  weighted  crucible  and  add  enough  powdered  sul- 


fur  to  cover  the  copper.  Place  the  cover  on  the  crucible  and  heat  gently  (with 
a  small  flame)  until  the  sulfur  ceases  to  burn  at  the  edges  of  the  cover,  but 
do  not  remove  the  cover  while  the  crucible  is  hot.  Then  heat  more  strongly 
until  the  bottom  of  the  crucible  just  becomes  dull  red.  Again  allow  to  cool 
about  15  minutes  and  weigh. 

Carefully  remove  the  cover  and  note  the  appearance  of  the  contents  of 
the  crucible,  but  do  not  touch  the  substance.  If  there  is  any  free  sulfur  on  the 
cover  or  wall  of  the  crucible,  replace  the  cover,  heat  the  crucible  and  cover, 
and  weigh  again.  In  order  to  check  the  final  weight,  add  a  small  quantity  of 
sulfur  and  repeat  the  experiment.  (While  waiting  for  the  crucible  to  cool 
answer  the  questions  in  the  following  paragraph.)  At  the  end  of  the  experi- 
ment remove  the  .substance  formed,  break  it,  and  note  the  differences  between 
its  properties  and  those  of  copper  and  sulfur.  To  check  the  weight  of  the 
empty  crucible,  clean  the  crucible  by  placing  it  in  a  porcelain  dish  containing 
a  little  .nitric  acid  and  warming  the  mixture,  rinse  the  crucible  with  distilled 
water,  heat  it,  let  it  cool,  andjweigh  i^_again.  By  means  of  a  table  summarize 
and  compare  the  results  obtained  in  both  parts  of  the  experiment.  If  there  is 
any  discrepancy  (beyond  10  mg.)  suggest  an  explanation. 

What  does  the  increase  in  weight  represent?  Is  any  free  sulfur  left  in  the 
crucible?  (Give  the  reasons  for  your  answers.)  What  would  the  increase 
in  weight  have  been  if  one  gram-atom  of  copper  had  been  used  in  the  experi- 
ment? How  does  this  weight  compare  with  the  atomic  weight  of  sulfur? 
What,  then,  is  the  formula  of  the  sulfide  of  copper  which  has  been  formed? 
Write  the  reaction  which  has  taken  place.  What  is  the  molecular  weight  of 
the  substance  formed? 

Problems.  I.  (a)  From  your  experimental  results  alone  calculate  the 
percentage  composition  of  the  copper  sulfide  formed,  (b)  From  the  atomic 
weights  of  copper  and  sulfur  and  the  molecular  weight  of  the  sulfide  calculate 
the  percentage  composition  of  pure  copper  sulfide.  Compare  the  results. 

2.  The  formula  of  hydrogen  sulfide  is  H,S.     How  much  sulfur  will  com- 
bine with  1.008  grams  of  hydrogen -to  form  this  compound?     What  weight  of 
hydrogen  sulfide  will  be  formed? 

3.  What  weight  of  iron  will  combine  with  32.07  grams  of  sulfur  to  form 
ferrous  sulfide,  FeS?     Write  the  reaction.     What  is  the  percentage  composition 
of  ferrous  sulfide?     How  much  ferrous  sulfide  would  be  needed  to  make  34-09 
grams  of  hydrogen  sulfide? 

ASSIGNMENT    III. 
THE  WEIGHT  OF  A  LITER  OF  OXYGEN. 

(Two  students  work  together.)  (Enter  the  date  and  title  of  this  experi- 
ment on  a  new  page.) 

When  solid  potassium  chlorate  is  strongly  heated  it  decomposes  with 
evolution  of  oxygen.  The  purpose  of  this  experiment  is  to  measure  the  volume 
of  a  definite  weight  of  oxygen  (obtained  from  potassium  chlorate).  This  may 
be  accomplished  by  heating  potassium  chlorate  in  a  hard-glass  test-tube,  and 
permitting  the  oxygen  evolved  to  displace  an  amount  of  water  equal  to  its  own 
volume.  The  loss  in  weight  of  the  potassium  chlorate  gives  the  weight  of  the 
oxygen  evolved,  and  the  volume  of  water  displaced  gives  the  volume  of  this 
amount  of  oxygen.  From  these  data  the  weight  of  i  liter  of  oxygen  can  be 
calculated. 

The  potassium  chlorate  decomposes  readily,  and  at  a  much  lower 
temperature  when  a  small  quantity  of  manganese  dioxide  is  present.  The  hard- 
glass  test-tube  may  then  be  replaced  by  a  heavy-walled  test-tube  of  ordinary 
easily-fusible  glass;  but  care  must  be  taken  not  to  heat  the  latter  tube  to  a 
higher  temperature  than  is  necessary  for  the  reaction.  The  manganese 
dioxide  is  a  "catalyzer"  in  this  reaction,  and  all  of  it  may  be  recovered  at 

7 


the  end  of  the  experiment  after  the  potassium  chlorate  has  been  decomposed 
into  potassium  chloride,  KC1,  and  oxygen. 

Two  students  working  together  obtain  from  the  office  a  heavy-walled  glass 
test-tube,  a  rubber  stopper,  glass  tube,  rubber  tube,  pinch  cock  and  clamp, 
One  student  signs  a  white  slip  for  "special  apparatus  for  Assignment  III"  , 
but  should  make  sure  that  the  other  shares  the  expense  with  him  if  any  portion 
of  this  apparatus  is  broken  or  damaged.  The  apparatus  should  be  returned  as  soon 
as  the  experiment  is  finished.  Each  student  must  keep  a  complete  record  of  the 
experiment  in  his  notebook. 

Experiment.  Set  up  the  apparatus  according  to  the  diagram  on  the  black- 
board and  follow  the  directions  given  by  the  instructor.  Plan  a  method  of  test- 
ing if  the  apparatus  is  air-tight,  and  be  sure  that  the  pressure  of  the  air  inside 
the  tube  at  the  beginning  of  the  experiment  is  the  same  as  the  atmospheric 
pressure  outside.  Do  not  begin  to  heat  the  tube  containing  the  potassium  chlo- 
rate until  the  instructor  has  examined  your  apparatus. 

Warm  the  heavy-walled  test-tube  carefully  to  dry  it,  holding  its  mouth 
lower  than  its  end  so  that  a  drop  of  water  may  not  run  down  the  side  on  to  the 
hot  glass  and  crack  the  tube.  Transfer  at  least  i  gram  of  powdered  potassium 
chlorate  (half  the  quantity  contained  in  the  small  tube)  to  a  porcelain  dish  and 
dry  it  by  holding  it  high  above  a  bunsen  flame  for  several  minutes ;  be  careful 
not  to  heat  the  substance  too  hot.  Set  the  chlorate  aside  to  cool,  weigh  the  test- 
tube  to  o.i  gram,  transfer  the  potassium  chlorate  to  the  test-tube,  and  weigh  to 
o.oi  g.  Add  a  small  quantity  of  manganese  dioxide  (not  more  than  50  mg.),  mix 
it  with  the  chlorate  by  jarring  the  tube,  remove  any  powder  from  the  outside  of 
the  tube  and  from  the  inside  near  its  mouth,  and  weigh  the  tube  carefully. 

Begin  the  experiment  by  opening  the  pinchcock  on  the  siphon  tube,  and 
heating  the  tube  containing  the  potassium  chlorate  very  gently  with  a  small 
flame.  Gradually  heat  the  tube  more  strongly  until  the  chlorate  melts  and  gas 
evolution  begins.  Then  continue  to  heat  the  tube  carefully  and  not  too  strongly 
until  enough  chlorate  has  decomposed  to  force  over  250  cc..  to  300  cc.  of  water 
into  the  receiving  beaker.  Now  discontinue  the  heating  and  allow  the  tube  to 
cool  to  room  temperature.  By  raising  the  beaker,  which  holds  the  displaced 
water,  equalize  the  level  of  the  water  in  the  beaker  and  the  flask,  and  then  close 
the  pinchcock  on  the  siphon  tube.  By  means  of  a  250  cc.  graduated  cylinder 
measure  the  amount  of  water  which  the  oxygen  has  forced  out  of  the  flask. 
Finally  weigh  carefully  the  hard  glass  tube  containing  the  partially  decomposed 
chlorate.  Question.  Why  is  it  necessary  to  cool  the  test-tube  and  to  have  the 
water  in  the  beaker  and  flask  at  the  same  level  before  closing  the  pinch-cock? 

Clean  the  test-tube  by  placing  water  in  it  and  shaking  the  mixture.  The 
manganese  dioxide  is  difficultly  soluble  in  water,  while  both  potassium  chlorate 
and  potassium  chloride  dissolve  readily.  Suggest  a  method  of  recovering  the 
manganese  dioxide  and  of  obtaining  a  mixture  of  dry  potassium  chloride  and 
potassium  chlorate  practically  free  from  manganese  dioxide. 

To  make  the  calculations  it  will  be  necessary  to  know  the  barometric 
pressure  at  the  time  you  perform  the  experiment,  and  the  temperature  of  the 
water  in  the  flask.  The  temperature  of  the  water  may  be  assumed  to  be  that 
of  the  room,  and  the  barometric  pressure  will  be  posted  on  the  blackboard. 
Enter  these  data  in  your  notebook  before  leaving  the  laboratory.  Below .  is 
given  a  table  of  the  vapor-pressure  of  water  at  different  temperatures. 

Vapor  Pressure    of    Water. 

Temp  °C.  Vapor  Pressure  Temp.  °C.  Vapor  Pressure 

14  2.2  cm.  mercury  24  1.2  cm.  mercury 

1 6  2.5     "  26  1.3 

18  2.8     "  28  1.5 

20  3-2  30  1-7 

22  3.5      '"  32  2.0 


:£-• , 

Questions  Assuming  that  the  levels  of  the  water  in  the  beaker  and  the 
flask  were  the  same  when  the  pinch-cock  was  closed,  what  was  the  total  pres- 
sure of  the  gas  in  the  flask?  What  was  the  partial  pressure  of  the  water-vapor? 
Of  the  oxygen? 

Calculate  the  volume  of  the  oxygen  at  i  atmosphere  pressure  and  o°  C, 
the  density  of  the  gas  under  standard  conditions,  and  the  weight  of  a  liter  of  the 
gas.  Compare  your  value  of  the  density  with  the  one  given  by  Cady,  page  23. 

Problems,  i.  From  your  experiment  calculate  the  weight  of  one  liter  of 
oxygen  at  2O°C  and  754  mm.  pressure. 

2.  Under    standard    conditions    wrhat    weight    of    oxygen    would    occupy 
22.4  liters?     (Use  the  density  obtained  in  the  experiment.)     How  is  this  num- 
ber related  to  the  atomic  weight  of  the  oxygen? 

3.  What  weight  of  oxygen  could  be  obtained  from   i   gram  of  potassium 
chlorate    (KC1O,)    by   completely   decomposing   it? 

ASSIGNMENT  IV. 
THE  REACTION   BETWEEN   ZINC  AND   SULFURIC  ACID. 

The  purpose  of  Assignment  IV  is  to  determine  the  quantity  of  zinc,  which, 
by  reacting  with  an  acid,  will  set  free  one  gram  atom  of  hydrogen.  A  weighed 
amount  of  zinc  is  allowed  to  react  completely  with  excess  of  sulfuric  acid  and 
the  hydrogen  liberated  is  collected  in  such  a  way  that  its  volume  can  be 
measured. 

Experiment.  Boil  about  500  cc.  water  in  your  largest  beaker  to  expel  the 
dissolved  air. 

Take  a  piece  of  zinc  weighing  about  half  a  gram.  If  necessary,  clean  it 
with  sand  paper.  Weigh  to  5  milligrams.  Do  not  try  to  cut  or  file  the  zinc  in 
order  to  obtain  any  definite  weight. 

Select  a  beaker  of  such  size  that  a  small  funnel  when  placed  in  it  can  be 
completely  covered  with  water.  Place  the  weighed  zinc  in  the  beaker,  place  the 
inverted  funnel  over  it,  and  pour  boiled  water  into  the  beaker  until  the  funnel 
is  completely  covered. 

Pour  boiled  water  into  a  quarter  liter  flask  until  the  water  completely  fills 
the  flask.  Without  waiting  for  the  water  to  cool,  moisten  a  piece  of  filter  paper 
slightly  larger  than  the  mouth  of  the  flask,  cover  the  mouth  of  the  flask  with1  the 
paper,  taking  care  that  no  bubble  of  air  remains  below  the  paper.  Invert  the 
flask  (over  an  empty  vessel)  and  lower  it  into  the  beaker  in  such  a  manner 
that  the  stem  of  the  funnel  enters  the  neck  of  the  flask.  If  a  bubble  of  air 
enters  the  flask  repeat  this  operation.  The  apparatus  now  consists  of  a  beaker 
containing  a  funnel  inverted  over  the  zinc,  and  a  flesk  filled  with  water  and 
inverted  over  the  funnel.  Place  this  apparatus  in  a  basin,  or  other  vessel,  to 
prevent  the  water  from  overflowing  on  the  desk  during  the  remainder  of  the 
experiment. 

Pour  into  the  water  5  cc.  to  8  cc.  concentrated  sulfuric  acid  (CAUTION*) 
If  the  action  of  the  acid  on  the  zinc  is  slow  add  about  5  cc.  of  the  laboratory 
solution  of  hydrochloric  acid;  stir  the  mixture  gently  but  be  careful  not  to  dis- 
place the  zinc  from  below  the  funnel.  The  hydrogen  evolved  rises  through  the 
funnel  into  the  flask,  but  should  not  displace  all  the  water  in  the  flask. 

When  the  zinc  has  all  dissolved  (except  a  few  dark-colored  flakes  of  im- 
purities of  negligible  weight),  place  the  apparatus  in  a  large  basin  of  water  and 
carefully  remove  the  beaker  and  funnel  without  allowing  any  air  to  enter  the 
inverted  flask.  Keep  the  flask  in  water  until  it  has  cooled  to  the  temperature  of 
the  water.  Then  raise  or  lower  the  flask  until  the  level  inside  and  outside  the 

*  Concentrated  sulfuric  acid  produces  dangerous  -burns  and  should  not  be  used  care- 
lessly. The  bottle  niust  not  be  removed  from  the  lead  tray.  Carefully  pour  just  enough 
acid  for  your  experiment  into  a  small  beaker. 


flask  is  the  same.  What  is  now  the  pressure  of  the  gas  inside  the  flask?)  While 
the  flask  is  in  this  position  cover  the  mouth  of  the  flask  with  the  palm  of  the 
hand,  remove  the  flask  from  the  water  and  invert  it.  While  the  gas  is  escaping 
test  to  prove  that  it  is  hydrogen. 

Fill  the  flask  by  means  of  a  graduated  cylinder,  noting  the  volume  of  water 
needed  to  take  the  place  of  the  gas  which  has  escaped.  Record  in  your  note- 
book the  barometric  pressure  (written  on  the  blackboard)  and  the  temperature 
of  the  water  in  which  the  flask  was  immersed. 

You  now  have  the  weight  of  zinc  taken,  and  the  volume,  at  a  definite  temp- 
erature and  pressure,  of  a  corresponding  amount  of  hydrogen  saturated  with 
water  vapor.  What  is  the  partial  pressure  of  the  water  vapor  at  the  tempera- 
ture of  the  experiment?  What  was  the  partial  pressure  of  the  hydrogen  in  the 
flask? 

Calculate  from  these  data : 

1.  The  volume  of  the  hydrogen  under  standard  conditions  of  temperature 
and  pressure. 

2.  The   weight   of   the  hydrogen.   .(The   density  of  pure  hydrogen   under 
standard  conditions  is  0.0000899.) 

3.  The  weight  of  zinc  that  would  have  liberated  i  gram  atom  of  hydro- 
gen.   Compare  this  number  with  atomic  weight  of  zinc. 

From  these  conclusions  write  the  equation  for  the  reaction.  The  formula 
of  sulfuric  acid  is  H2SO4,  and  both  hydrogens  are  replaced  by  zinc. 

Problems,  i.  (a)  From  the  density  of  hydrogen  calculate  the  weight  of 
22.4  liters  of  this  gas  under  standard  conditions.  What  is  the  true  molecular 
weight  of  hydrogen?  What  is  the  formula  of  this  gas?  Have  you  written  the 
above  equation  correctly? 

(b)  If  you  had  not  been  given  the  density  of  hydrogen  in  the  above  ex- 
periment, how  could  you  have  calculated  a  value  from  the  molecular  weight? 
What  percentage  error  would  you  have  made? 

2.  (a)     What  weight  of  pure   zinc   sulfate    (ZnSO4)    could  be   obtained 
from  the  weight  of  zinc  used? 

(b)  What  is  the  minimum  weight  of  sulfuric  acid  (H2SO4)  necessary  in 
the  experiment? 

3.  If  hydrochloric  acid  (HC1)  had  been  used  instead  of  sulfuric  acid,  how 
much  hydrogen  would  have  been  liberated?     Write  the  equation.     What  weight 
of  solid  zinc  chloride  could  have  been  obtained? 

Note  on  Glass  Manipulation. 

To  bend  a  piece  of  ordinary  glass  tubing,  hold  it  with  both  hands  in  a  fan- 
shaped  gas  flame  and  rotate  it  slowly  between  the  thumb  and  fingers  until  a 
2.y2  to  3  inch  portion  is  uniformly  heated  and  is  soft  enough  to  be  bent  to  the 
proper  angle.  Set  it  aside  to  cool;  glass  will  remain  hot  enough  to  burn  the 
hand  for  some  time  after.it  no  longer  appears  to  be  hot. 

To  cut  glass  tubing,  scratch  it  with  a  file  at  the  proper  place,  grasp  it  firmly 
on  each  side  of  this  mark  (protecting  the  hands  with  a  cloth),  and  bend  the 
tube  away  from  the  mark. 

Always  remove  the  sharp  edges  of  freshly  cut  glass  at  once  with  a  file,  or 
by  heating  in  a  gas  flame. 

To  draw  down  a  piece  of  tubing  to  a  capillary,  heat  a  portion  about  i 
inch  long  in  an  ordinary  gas  flame  to  a  higher  temperature  than  was  necessary 
in  bending  the  tubing.  Hold  the  tube  with  both  hands  and  rotate  it  to  ensure 
uniform  heating  and  prevent  the  hot  portion  from  sagging.  Withdraw  from  the 
flame  and  draw  apart  slowly  to  obtain  a  thick-walled  capillary.  Mount  your 
piece  of  platinum  wire  (first  used  in  Assignment  VI)  by  inserting  one  end  of 
the  wire  in  the  capillary  near  the  tube,  and  heating  until  the  glass  closes  firmly 
around  the  wire. 

10 


ASSIGNMENT   V. 
THE  ANALYSIS  OF  COPPER  OXIDE 
(Two  students  work  together.) 

The  purpose  of  Assignment  V  is  to  determine  the  relative  weights  of 
copper  and  oxygen  in  copper  oxide.  Hydrogen  gas  reacts  with  hot  copper 
oxide  to  form  metallic  copper  and  steam.  Two  measurements  are  necessary: 
the  weight  of  the  copper  oxide  used,  and  the  loss  in  weight  when  it  is  com- 
pletely reduced  to  copper.  Question :  What  does  this  loss  of  weight  repre- 
sent? 

One  of  the  students  signs  a  white  slip  for  "special  apparatus  for  Assign- 
ment V",  consisting  of  a  thistle  tube,  a  clamp,  calcium  chloride  tube,  with  two 
rubber  stoppers,  2  glass  tubes,  2  rubber  tubes,  and  2  short  glass  rods,  and 
returns  the  apparatus  as  soon  as  the  experiment  is  finished. 

Out  of  a  wash  bottle  make  a  hydrogen  generator  by  fitting  it  with  a 
thistle  tube  (extending  almost  to  the  bottorri  of  the  flask)  and  an  outlet  tube 
bent  at  right  angles  (see  note  preceding  this  Assignment).  Place  in  the  wash 
bottle  about  10  grams  of  zinc,  and  (to  ensure  a  rapid  reaction  between  Zn 
and  H2SO4)  cover  with  a  very  dilute  solution  of  copper  sulfate  made  by 
diluting  10  cc.  of  the  laboratory  solution  to  100  cc.  with  water.  To  the  out- 
let tube  attach  a  "drying  tube"  (containing  solid  calcium  chloride,  which  has 
the  property  of  absorbing  moisture).  Make  sure  that  the  apparatus  is  air- 
tight, and  wrap  the  flask  in  a  towel. 

Set  up  the  remainder  of  the  apparatus  according  to  the  directions  of  the 
instructor.  Dry  the  thick-walled  glass  tube  that  is  to  contain  the  copper 
oxide  by  heating  it  gently.  When  it  is  cool  weigh  it  carefully,  together  with 
any  portion  of  the  apparatus  that  may  come  in  contact  with  the 
copper  oxide.  Place  in  the  tube  about  i  gram  of  copper  oxide, 
wipe  off  any  particles  that  are  not  in  the  portion  of  the  tube 
that  is  to  be  heated.  Weigh  again  carefully  to  obtain  the  weight  of  copper 
oxide  used.  Attach  the  apparatus  to  the  hydrogen  generator,  pour  10  cc.  con- 
centrated sulfuric  acid  (Caution!)  down  the  thistle  tube,  and  allow  the  hydro- 
gen to  pass  through  the  apparatus  until  it  has  swept  out  the  air.  (Caution! 
Do  not  place  a  flame  near  the  outlet  nor  heat  the  oxide  while  the  apparatus 
contains  a  mixture  of  oxygen  and  hydrogen.  A  dangerous  explosion  might 
result.)  Suggest  a  simple  test  to  determine  when  the  hydrogen  is  no  longer 
mixed  with  oxygen. 

When  pure  hydrogen  is  passing  over  the  copper  oxide,  begin  to  heat  it 
very  gently  with  a  small  flame  and  continue  to  heat  cautiously  until  all  the 
oxide  changes  color.  If  moisture  collects  in  the  farther  end -of  the  tube  drive 
it  out  by  heating  the  tube  carefully.  Question :  Where  does  this  moisture 
come  from  ? 

Allow  the  tube  to  cool  in  the  current  of  hydrogen,  and  weigh  it.  If  you 
have  time,  check  this  result  at  once  by  repeating  the  heating  and  weighing;  if 
not,  set  the  tube  aside  in  order  that  you  may  do  so  if  the  results  of  the  following 
calculations  are  unsatisfactory. 

From  the  results  of  your  experiments  calculate  the  percent  of  copper 
and  oxygen  in  copper  oxide. 

How  many  grams  of  copper  are  combined  with  i  gram  atom  of  oxygen? 
Compare  this  figure  with  the  atomic  weight  of  copper.  What,  then,  is  the 
formula  of  this  oxide  of  copper?  Write  the  equation  for  the  reaction.  What 
is  the  true  percent  of  copper  and  oxygen  in  this  oxide  of  copper? 

Problems,  i.  (a)  Another  oxide  of  copper  is  known  which  contains 
88.8%  of  copper.  What  is  its  formula? 

(b)     Which   of  the   oxides   is   cuprous  oxide   and   which   cupric   oxide? 

ii 


(c)  What   are   the   formulas   of   cuprous   and   cupric   sulfides? 

(d)  Two  of  the   oxides   of   iron   are   FeO   and   Fe2CX :   which   is   ferrous 
oxide  arid  which  is  ferric  oxide? 

2.  How    many   grams   of   water   can   be    formed    from    i    gram   of    cupric 
oxide?     How   many  grams   of  hydrogen  would  be  needed?     What   volume  of 
hydrogen  measured  at  20°  C  and  i   atmosphere  would  be  needed? 

3.  How  many  grams  of  zinc  would  be  needed  to  generate  enough  hydro- 
gen   (by   reacting  with   an   acid)    to   reduce   completely  one  gram   molecule  of 
cupric  oxide?     Compare  this  weight  with  the  atomic  weight  of  zinc. 

ASSIGNMENT  VI. 
PROPERTIES  OF  AQUEOUS   SOLUTIONS  OF  ACIDS  AND  BASES. 

(a)  Acids.  Prepare  a  solution  of  each  of  the  common  acids,  hydro- 
chloric, nitric  and  sulfuric,  for  use  in  the  following  experiments,  by  adding 
50  cc.  water  to  5  cc.  of  the  "dilute"  laboratory  solution  of  each  of  these  acids. 
Label  each  solution. 

Taste.  Taste  each  solution  by  dipping  a  glass  rod  into  the  liquid  and 
touching  it  to  the  tongue.  (Caution:  Do  not  taste  any  substance  in  the  labor- 
atory unless  directed  to  do  so.) 

Action  on  Indicators.  To  about  20  cc.  of  distilled  water  in  a  beaker  add 
enough  blue  litmus  solution  to  give  a  distinct  but  not  strong  color.  Divide 
this  solution  into  three  portions  and  to  each  add  a  few  drops  of  a  different  acid. 

Repeat,  using  methyl  orange  instead  of  litmus. 

Withdraw  a  drop  of  any  one  of  the  acids  on  a  glass  rod  and  touch  it  to 
blue  litmus  paper. 

Repeat,  using  red  litmus  paper. 

Action  on  Sodium  Carbonate.  To  5  cc.  of  each  acid  in  separate  test 
tubes  add  a  little  powered  sodium  carbonate. 

Action  on  Metallic  Zinc.  To  5  cc.  of  each  acid  add  small  pieces  of  zinc. 
Is  the  same  gas  given  off  in  each  case?  Test  it.  Repeat  the  experiment  with 
zinc  and  the  ordinary  laboratory  solution  of  the  acids. 

Summarize  the  properties  common  to  the  acids  you  have  studied.  Note 
that  any  substance  is  recognized  by  its  properties.  Any  solution  ^  which  has 
the  properties  thus  summarized  is  an  acid  solution  and  contains  the  substance 
hydrogen  ion.  The  symbol  is  H+. 

In  addition  to  the  properties  of  hydrogen  ion,  the  solution  of  each  acid 
has  another  group  of  properties  peculiar  to  itself  by  means  of  which  the  acids 
may  be  distinguished  from  one  another.  Try  the  following  experiments : 

To  15  cc.  water  add  about  2  cc.  silver  nitrate  solution.  Divide  the  mix- 
ture into  three  portions  and  to  each  add  a  few  cc.  of  a  different  acid. 

In  a  similar  set  of  experiments  use  barium  chloride  solution  instead  of 
silver  nitrate. 

State  how  the  three  acids  may  be  distinguished. 

Thus,  in  addition  to  hydrogen  ion,  hydrochloric  acid  solution  contains 
chloride  ion,  C\~ ;  nitric  acid  solution  contains  nitrate  ion,  NO3~;  and  sulfuric 
acid  solution  contains  sulfate  ion,  SO4~~. 

Other  properties  of  ions  will  be  discussed  in  the  lectures,  and  additional 
evidence  of  the  existence  of  ions  in  aqueous  solutions  will  be  presented.  The 
study  of  ionization  will  be  continued  in  the  laboratory  in  Assignment  IX. 

The  chemical  formulas  of  these  three  substances,  whose  aqueous  solutions 
contain  hydrogen  ion,  are  : 

Hydrochloric  acid    HC1, 
Nitric  acid,  HNO3, 

Sulfuric  acid,  H2SO4. 

Calculate  the  molecular  weight  of  each  acid.     Pure  HC1  is  a  gas.  and   it 

12 


is  also  named  hydrogen  chloride  and  hydrochloric  acid  gas.  Pure  HNO3  and 
H2SO4  are  liquids. 

(b)  Bases.  To  5  cc.  portions  of  the  laboratory  solutions  of  sodium 
hydroxide  and  potassium  hydroxide  (taken  separately)  add  50  cc.  water.  Use 
these  two  solutions  and  the  undiluted  laboratory  solution  of  barium  hydroxide 
in  the  following  experiments : 

Taste.  Taste  each  of.  these  three  solutions  by  withdrawing  a  drop  on  a 
glass  rod  and  touching  it  to  the  tongue. 

Touch.     Rub  a  drop  of  each  solution  between  the  thumb  and  finger. 

Action  on  Indicators.  Treat  drops  of  each  of  the  three  solutions  with  red 
and  with  blue  litmus  paper. 

To  5  cc-  of  each  of  the  same  solutions  add  a  drop  of  litmus  solution.  To 
the  sodium  hydroxide  solution  containing  the  indicator  add  hydrochloric  acid 
slowly  and  note  the  change  of  color.  Then  add  sodium  hydroxide  solution 
until  a  change  is  noticed. 

Repeat  the  last  experiment,  using  fresh  5  cc.  portions  of  the  same  hydrox- 
ide solutions  and  a  solution  of  the  indicator  methyl  orange. 

Repeat,   using  a  solution  of  the  indicator  phenolphthalein. 

Action  with  Ferric  Chlcride  Solution.  To  5  cc.  portions  of  each  of  the 
hydroxide  solutions  add  ferric  chloride  drop  by  drop. 

Summarize  the  properties  common  to  the  solutions  of  the  bases  you  have 
studied.  Each  solution  contains  the  substance  hydroxide  ion,  OH~. 

In  addition  to  hydroxide  ion  the  solution  of  each  base  also  has  properties 
characteristic  of  the  base. 

To  5  cc.  of  each  hydroxide  solution  add  a  few  drops  of  sulfuric  acid. 

Test  a  drop  of  each  of  the  hyroxide  solutions  in  a  colorless  gas  flame  by 
means  of  a  looped  platinum  wire,  as  demonstrated  by  the  instructor.  Can  the 
flame  test  be  used  to  prove  the  presence  of  sodium? 

State  how  these  bases  may  be  distinguished. 

In  addition  to  hydroxide  ion,  sodium  hydroxide  solution  contains  sodium 
ion,  Na+ ;  potassium  hydroxide  contains  potassium  ion,  K+ ;  and  barium 
hydroxide  contains  barium  ion,  Ba++. 

Each  of  the  hydroxides  in  the  pure  state  is  a  solid.  The  chemical  form- 
ulas for  the  hydroxides  are: 

^Sodium  hydroxide,         NaOH, 
Potassium  hydroxide,    KOH, 
Barium  hydroxide,         Ba(OH)2. 
What   is  the  molecular  weight  of   each   hydroxide? 

Problems,  i.  Make  as  simple  a  table  as  you  can  to  show  the  colors 
obtained  with  different  indicators  in  acid  and  basic  solutions. 

2.  Four  unlabelled  beakers,  which  are  known  to  contain  solutions  of 
sulfuric  acid,  nitric  acid,  barium  hydroxide,  and  sodium  hydroxide,  are  to  be 
identified  as  quickly  as  possible.  State  what  experiments  you  would  perform, 
and  note  clearly  the  progress  towards  identification  that  would  be  made  in 
each  experiment. 

ASSIGNMENT   VII. 
REACTIONS  BETWEEN  ACIDS  AND  BASES. 

Experiment.  Evaporate ,  to  dryness  in  a  porcelain  dish  a  few  drops  of 
(a)  hydrochloric  acid  solution,  (b)  sodium  hydroxide  solution.  In  each  case 
note  if  there  is  a  residue,  add  a  little  water,  and  test  the  solution  with  an 
indicator.  Questions:  Taking  into  account  that  HC1  and  NaOH  are  both 
stable  substances,  state  what  gases  were  given  off  when  the  solutions  were 
evaporated.  What  conclusion  do  you  draw  in  regard  to  the  volatility  of 
hydrogen  chloride  and  sodium  hydroxide?  How  is  sodium  hydroxide  pre- 
13 


pared  from  metallic  sodium  and  water,  hydrogen  chloride  from  hydrogen  and 
chlorine  gases? 

Experiment.  To  10  cc.  sodium  hydroxide  solution  in  a  beaker  add  hydro- 
chloric acid  solution  slowly  until  a  drop  of  the  mixture,  after  stirring,  reacts 
acid  to  litmus  paper.  Evaporate  in  a  porcelain  casserole.  While  the  solution 
is  boiling,  test  again  with  litmus  paper,  and  if  the  solution  is  no  longer  acid 
add  HC1.  When  the  solid  begins  to  separate,  use  a  small  flame,  and,  if  neces- 
sary, withdraw  the  flame  and  stir  the  mixture.  Finally  heat  until  the  mixture 
is  thoroughly  dry.  Allow  the  dish  to  cool,  and  note  the  appearance  of  the 
residue.  Dissolve  it  in  about  10  cc.  of  water.  Test  the  solution  with  blue 
litmus  paper  and  with  phenolphthalein  test  a  drop  on  a  looped  platinum  wire 
in  a  colorless  gas  flame  and  test  a  small  portion  of  the  solution  with  silver 
nitrate  in  the  presence  of  nitric  acid.  Questions.  Does  this  solution  have 
either  acid  or  basic  properties?  What  does  the  solution  contain?  What  was 
the  solid  substance  obtained  on  evaporation?  What  substances  were  vola- 
tilized? Why  was  hydrochloric  acid  added  in  excess  instead  of  sodium 
hydroxide  ? 

Write  the  equation  (non-ionic)  for  the  reaction  between  sodium  hydrox- 
ide and  hydrochloric  acid,  and  state  what  has  become  of  the  hydrogen  of  the 
acid  and  the  hydroxide  radical  of  the  base.  t 

This  reaction  is  typical  of  the  action  of  any  acid  on  any  base. 

Write  the  equations  (non-ionic)  for  the  following  reactions,  consulting 
a  text-book  if  necessary. 

1.  Nitric  acid  and  sodim  hydroxide. 

2.  Sulfuric   acid    and   sodium   hydroxide. 

3.  Hydrochloric   acid   and   barium   hydroxide. 

Questions  and  Problems.  i.  Calculate  the  weight  of  sodium  chloride 
that  will  be  obtained  on  neutralizing  one  gram  of  sodium  hydroxide  with 
hydrochloric  acid.  How  many  molecules  of  HC1  react  with  one  molecule  of 
Ba(OH)2?  What  weight  of  barium  hydroxide  will  be  obtained  on  neutral- 
izing one  gram  of  barium  hydroxide  with  hydrochloric  acid? 

2.  What  weight  of  barium  hydroxide  will  react  with  one  mol  of  hydro- 
gen chloride?     This  is   called  one   equivalent  of  barium  hydroxide. 

Suggest  a  similar  definition  for  one  equivalent  of  sulfuriq  acid  from  its 
reaction  with  sodium  hydroxide.  Note  that  one  equivalent  of  any  acid  will 
react  with  one  equivalent  of  any  base  and  the  resulting  solution  will  contain 
one  equivalent  of  a  corresponding  salt. 

How  many  equivalents  of  barium  hydroxide  will  react  with  one  mol  of 
sulfuric  acid?  How  many  equivalents  of  barium  sulfate  will  be  formed? 

3.  Concentration  of  a  Solution.     The   concentration  of  a  solution  is   the 
amount  of  the  dissolved  substance  contained  in  a  unit  volume  of  the  solution. 

The  molecular  weight  in  grams  (the  mol)  is  frequently  adopted  as  a  unit 
of  mass,  and  the  liter  is  the  standard  unit  of  volume.  A  solution  which  con- 
tains in  one  liter  one  mol  of  the  dissolved  substance  is  called  a  mold  solution. 

How  many  grams  of  sulfuric  acid  are  contained  in  a  liter  of  a  molal 
solution? 

How  many  grams  of  sodium  chloride  are  contained  in  a  liter  of  a  one- 
tenth  molal  solution?  (o.i  M  NaCl  designates  this  solution.) 

4.  When  we  are  dealing  only  with  neutralization  reactions  it  is   conven- 
ient to  use  the  equivalent   (of  the  acid,  or  base,  or  salt)    as  the  unit  of  mass. 
A  solution  which  contains  one  equivalent  in  a  liter  is  called  a  normal  solution. 

How  many  grams  of  HC1  are  contained  in  a  liter  of  a  o.io  normal  solu- 
tion of  hydrochloric  acid  (o.io  N  HC1)  ? 

How  many  grams  of  Ba(OH)2  are  contained  in  a  liter  of  o.io  N  barium 
hydroxide  solution? 

How    many   cc.    of  o.io  N   HC1   must   be   added   to    100   cc.   of  o.io  N 


Ba(OH)2  to  produce  a  neutral  solution  of  barium  chloride?  What  color 
changes  would  be  observed  while  the  acid  is  being  added  to  the  barium 
hydroxide  solution  (i)  if  methyl  orange  were  present,  (2)  if  phenolphthalein 
weije  present? 

ASSIGNMENT  VIII. 

TlTRATION    OF    ACIDS    AND    BASES.       AN    ILLUSTRATION    OF 

VOLUMETRIC  ANALYSTS. 

.Ac-;  —  r^^nt  VTII  the  purpose  is  to  measure  the  volumes  of  solutions 
ofr,acid  and  base  which  will  exactly  neutralize  each  other.  This  neutral  point, 
called  the  end-point,  is  determined  by  means  of  some  indicator  like  litmus  or 
phenolphthalein.  The  whole  experiment  is  called  titration.  Question. — Does 
the  color  of  an  indicator  used  in  titration  change  gradually  or  suddenly  as  the 
solution  is  neutralized?  If  necessary,  perform  an  experiment  to  determine 
this  point.) 

Questions.  To  what  volume  would  you  dilute  10  cc.  of  a  6-normal 
solution  to  prepare  a  3-normal  solution?  How  would  you  prepare  a  o.5-normal 
solution  from  a  6-normal  solution? 

Experiment.  Prepare  about  400  cc.  approximately  o.5-normal  sodium 
hydroxide  solution  from  the  laboratory  6-normal  solution.  Place  this  0.5- 
normal  solution  in  your  500  cc.  flask.  Prepare  also  about  250  cc.  approxi- 
mately o.5-normal  sulfuric  acid  solution  from  the  laboratory  6-normal  solu- 
tion, and  place  it  in  another  flask.  Cork  and  label  each  flask.  Shake  the 
flasks  in  order  that  the  concentration  of  the  solutions  will  be  uniform  through- 
out. Clean  two  other  flasks,  rinse  them  with  distilled  water,  and  set  them 
aside  to  drain,  in  order  that  they  may  be  ready  for  use  later  in  the  experiment ; 
label  one  "known  H2SO4  solution",  and  the  other  "unknown  HC1  solution". 

Note.  The  volumes  of  the  solutions  used  in  a  titration  are  measured  by 
means  of  burettes.  These  are  glass  tubes,  marked  off  in  cubic  centimeters  and 
tenths  (or  fifths)  of  cubic  centimeters,  and  provided  with  a  glass  tip  out  of 
which  the  solution  may  be  drawn  by  opening  a  pinch-cock. 

Apply  at  the  office  for  two  burettes  and  clamps,  and  return  them  the  same 
day  before  the  office  closes.  Sign  a  yellow  temporary  order  slip  when  you 
receive  these  articles,  and  a  blue  return  slip  when  you  return  them. 

Experiment.  Fill  each  burette  with  distilled  water.  Note  that  air  may 
be  removed  from  the  small  tube  below  the  pinch-cock  by  tilting  the  tip  upward 
and  allowing  the  liquid  to  flow  through  the  pinch-cock.  Practice  reading  a 
burette :  bring  your  eye  to  the  same  level  as  the  liquid  and  note  the  reading 
of  the  burette  corresponding  to  the  bottom  of  the  meniscus ;  repeat  until  con- 
secutive readings  check  to  better  than  0.05  cc.  Question.  Why  is  it  import- 
ant to  have  the  eye  at  the  same  level  as  the  liquid  before  making  a  reading? 
Never  attempt  to  adjust  the  volume  of  the  solution  in  a  burette  so  that  the 
reading  will  be  some  exact  amount.  Allow  the  water  to  flow  slowly  out  of 
the  burette.  If  drops  remain  on  the  inner  surface  of  a  burette,  exchange  it  at 
the  office  for  a  clean  one. 

Rinse  one  burette  with  a  little  of  your  approximately  o.5-normal  H2SO4 
solution,  and  fill  the  burette  with  this  solution.  Rinse  and  fill  the  other  burette 
with  the  o.5-normal  NaOH  solution. 

Record  the  readings  of  the  burettes  side  by  side  in  your  note-book;  run 
about  10  cc.  of  the  acid  solution  into  a  clean  beaker  or  flask  standing  on  white 
paper,  and  record  the  final  burette  reading  under  the  initial  reading.  Add  two 
drops  of  phenolphthalein,  and  about  20  cc.  distilled  water.  Then  run  in  the 
sodium  hydroxide  solution  from  the  other  burette,  a  little  at  a  time,  and  to- 
wards the  end,  very  carefully,  a  drop  or  two  at  a  time,  stirring  the 
mixture  constantly,  until  the  faintest  perceptible  .permanent  pink  color  is 

15 


obtained.  Wash  down  the  inside  'of  the  beaker  by  means  of  a  jet  of  water 
from  the  wash  bottle.  If  too  much  of  the  basic  solution  is  added,  decolorize 
the  solution  by  adding  a  little  of  the  acid  solution  and  4etermine  the  end-point 
again.  Record  the  final  readings  of  each  burette,  and  the  actual  volumes  of 
each  solution  used  in  the  titration.  Calculate  the  ratio  of  the  two  concentra- 
tions and  state  which  is  the  more  concentrated  solution. 

Repeat  this  experiment,  using  about  15  cc.  and  then  about  20  cc.  of  the 
acid  solution,  and  in  each  case  make  the  same  calculations.  Do  not  fill  up  the 
burette  before  a  titration  unless  there  is  not  enough  solution  in  it  for  the 
titration. 

Compare  the  ratios  of  the  concentrations  calculated  from  these  three  titra- 
tions.  If  any  result  differs  from  the  average  by  more  than  i%,  perform 
additional  titrations  until  you  are  satisfied  that  you  have  determined  the  ratio 
of  the  concentrations  with  an  accuracy  better  than  i%. 

Take  your  two  clean,  dry,  labelled  flasks  to  the  office  to  obtain  a  sulfuric 
acid  solution  of  known  concentration,  and  a  hydrochloric  acid  solution  whose 
concentration  you  are  to  determine. 

Empty  your  acid  burette,  fill  it  with  the  sulfuric  acid  solution  of  known 
concentration,  and  determine  as  before  from  three  (or  more)  titrations  the 
relative  concentrations  of  this  solution  and  your  sodium  hydroxide  solution. 
Never  use  less  than  10  cc.  of  any  solution  in  a  titration.  Calculate  the  con- 
centration of  your  sodium  hydroxide  solution. 

Finally  titrate  the  unknown  hydrochloric  solution  against  the  same  sodium 
hydroxide  solution,  and  calculate  the  concentration  of  the  acetic  acid  solution. 
Report  this  result  at  once  to  your  instructor.  If  the  result  is  not  sufficiently 
accurate,  first  look  for  an  error  in  your  calculations,  but,  if  necessary,  repeat 
the  experiment. 

Make  a  list  of  the  sources  of  error.  (Many  of  them  have  been  mentioned 
in  the  above  directions.) 

Save  the  remainder  of  the  NaOH  solution, w7hose  concentration  you  have 
determined,  in  a  corked  flask  for  use  in  Assignment  X. 

Problems.  i.  If  the  error  in  measuring  out  a  volume  of  solution  by 
means  of  a  burette  is  o.io  cc.  what  is  the  percentage  error  if  i  cc.  of  solution 
is  measured?  If  20  cc.  of  solution  are  measured?  Why  is  it  important  not 
to  use  less  than  10  cc.  in  any  titration? 

2.  How  many  cubic  centimeters  of  i  molal  sulfuric  acid  will  contain  o.oi 
equivalent?     How  many  grams  of  sodium  hydroxide  would  be  needed  to  neu- 
tralize this  amount  of  acid? 

3.  Chemically    pure     ("C.  P.")     sulfuric    acid,    nitric    acid,    hydrochloric 
acid   and   ammonia,   as   supplied     by     the     manufacturers,      are      concentrated 
aqueous  solutions  of  these   substances.     The  concentration  of  each  solution  is 
guaranteed  not  to  be  less  than  a  certain  minimum  value,  and  this  is  tested  by 
measuring   the    density    (specific    gravity).      The    following   table    contains    the 
density  and  the  percentage  composition  by  weight  of  the  concentrated  labora- 
tory  reagents. 

Density  %  by  weight         Normal   Concentration 

H2SO4  1.84  95-6%  H.SO, 

HNO3  1.42  69.8%  HNO3 

HC1  1.19  37.2%  HC1 

NH4OH  0.90  28.3%  NH3 

For  each  reagent  calculate  the  normal  concentration  and  record  the 
results  in  the  fourth  column  of  the  table. 

CAUTION!  The  concentrated  acids,  especially  sulfuric  and  nitric,  pro- 
duce dangerous  burns  and  should  not  be  used  carelessly.  These  reagents  are 
kept  on  lead  shelves  and  must  not  be  used  unless  special  directions  are  given. 

16 


ASSIGNMENT  IX. 
STRONG  ACIDS.     IONIC  EQUATIONS. 

In  Assignments  VI  and  VII  the  general  properties  of  acids  and  bases 
were  studied,  and  it  was  recognized  that  acid  solutions  contain  hydrogen  ion, 
and  basic  solutions  hydroxide  ion.  The  purpose  of  Assignments  IX  and  X 
is  to  compare  the  concentrations  of  hydrogen  ion  in  solutions  of  different 
acids  and  to  introduce  the  writing  of  ionic  equations. 

Experiment.  Prepare  a  one-normal  solution  (about  60  cc.)  hydrochloric 
acid  from  the  6  TV  laboratory  reagent  (see  questions  below)  and  from  this 
solution  prepare  about  50  cc.  each  of  o.io  N,  .01  N  and  .001  N  solutions.  Be 
careful  to  shake  each  solution  after  preparing  it  from  a  more  concentrated 
solution  and  water. 

Questions.  What  volume  of  6  N  acid  is  needed  to  make  100  cc.  of  i  AT 
acid? 

How  much  water  should  be  added  to  10  cc.  of  6  N  acid  to  make  i  N  acid? 

How  much  water  should  be  added  to  10  cc.  of  i  N  acid  to  make  o.io 
N  acid? 

What  volume  of  o.oi  N  acid  can  be  made  from  5  cc  of  o.i  acid? 

If  hydrochloric  acid  is  ionized  completely  what  is  the  concentration  of 
hydrogen  ion  and  of  chloride  ion  in  each  of  the  solutions  prepared? 

Experiment.  Pour  into  marked  test  tubes  10  cc.  of  each  solution  (N, 
o.i  N,  o.oi  TV  o.ooi  N},  and  pour  10  cc.  of  wrater  into  a  fifth  test  tube.  Add 
to  each  solution  from  a  glass  tube  a  single  drop  of  methyl  violet  solution. 
Hold  the  tubes  in  a  vertical  position,  look  down  at  the  surface  of  the  solutions, 
record  the  color  of  each  solution,  and  note  the  smallest  concentration  of 
hydrochloric  acid  that  shows  with  this  indicator  a  different  color  from  water. 

State  how  the  indicator,  methyl  violet,  may  be  used  to  determine  the  ap- 
proximate concentration  of  a  hydrochloric  acid  solution.  (The  color  in  the  more 
concentrated  solutions  will  fade  on  standing.  It  may  be  restored  by  adding 
another  drop  of  the  indicator.)  Set  the  o.oi  TV  solution  aside  for  use  in  a  later 
experiment  (Assignment  XIII). 

Repeat  the  experiment  with  nitric  acid  and  with  sulfuric  acid,  starting  in 
each  case  with  the  6  TV  laboratory  acid.  Compare  the  colors  obtained  with 
the  different  acids.  Questions.  Are  the  colors  characteristic  for  each  acid? 
If  not,  what  substance  determines  the  color?  If  hydrochloric  acid  is  com- 
pletely ionized  what  conclusion  can  you  draw  with  respect  to  the  ionization  of 
nitric  acid  and  sulfuric  acid?  Do  these  experiments  prove  that  these  three 
acids  are  highly  ionized? 

Summarize  briefly  the  evidence  presented  in  the  lectures  that  these  acids, 
the  bases  sodium  hydroxide  and  potassium  hydroxide,  and  nearly  all  salts  are 
highly  ionized  in  dilute  solutions.  Such  substances  are  called  strong  electro- 
lytes, these  three  acids  are  strong  acids,  and  sodium  hydroxide  is  a  strong  base. 
The  term  strong  salt  is  seldom  used,  because  there  are  so  few  slightly  ionized 
salts.  In  pure  water  the  concentration  of  H+  and  OH~  ions  is  extremely 
small  and  it  is  an  example  of  a  very  weak  electrolyte. 

We  shall  next  consider  what  ionic  reaction  takes  place  when  a  strong 
acid  neutralizes  a  strong  base.  The  equation, 

H,SO4   +  2NaOH  =  2H2O   +   Na,SO4, 

which  we  have  been  using. hitherto,  states  that  i  mol  or  2  equivalents  of  sul- 
furic acid  and  2  mols  or  2  equivalents  of  sodium  hydroxide  react,  and  that  the 
resulting  solution  contains  i  mol  or  2  equivalents  of  the  salt  sodium  sujfate,  and 
2  more  mols  of  water  than  were  formerly  present ;  but  it  does  not  show  the  differ- 
ence between  the  ions  present  in  the  two  initial  solutions  and  in  the  final  solution. 
This  difference  may  be  stated  in  words,  or  expressed  briefly  in  the  following 
equation : 

17 


+  SO4-)   +   (2Na+  4-  2OH-)  =  2H2O  +  2Na+  +  SO4~. 

The  Na+  and  SO4™  are  the  same  in  the  initial  and  final  solutions,  and  it  is  not 
correct  to  say  that  they  have  united  to  form  Na2SO4.  The  reaction  that  has 
taken  place  is 

H+  +  OH-  =  H2O. 

This  equation  states  that  i  mol  or  i  equivalent  of  hydrogen  ion  has  combined 
with  i  mol  or  i  equivalent  of  hydroxide  ion  to  form  i  mol  of  water.  Write 
the  ionic  equations  for  the  neutralization  of  hydrochloric  acid  and  nitric  acid 
solutions  by  sodium  hydroxide  solution.  State  briefly  the  evidence  presented 
in  the  lectures  to  show  that  the  reaction  between  any  strong  acid  and  any 
strong  base  is  the  same  in  all  cases.  What  reaction  will  take  place  whenever 
a  solution  with  acid  properties  is  mixed  with  a  solution  that  has  basic  prop- 
erties? 

Problems. — i.     Write   ionic   equations   for  the   following   reactions: 

(a)  A  sodium  chloride  solution  is  evaporated  to  dryness    (solids  are 
not  ionized). 

(b)  A   sodium   sulphate   solution   is  evaporated   to  dryness. 

(c)  Hydrogen    chloride    gas    is    dissolved    in    water    (gases    are    not 

ionized). 

(d)  A    precipitate    of    silver    chloride,    AgCl,    is    formed    by    mixing 

solutions  of  silver  nitrate  and  sodium  chloride. 

(e)  A   precipitate    of    barium    sulfate,    BaSO4,    is    formed   by   mixing 

solutions   of   barium   chloride   and   sodium   sulfate. 

2.  Does  anything  happen  in  the  following  experiments? 

(a)  Dilute    solutions    of    sodium    chloride    and   potassium   nitrate    are 

mixed. 

(b)  Dilute  solutions  of  sodium  chloride  and  nitric  acid  are  mixed. 

3.  What  is  the   concentration  of   Na+,   of   Cl~,   in  a  solution  prepared  by 
dissolving  1.17  grams  sodium  chloride  in   i   liter  of  water? 

ASSIGNMENT  X. 
WEAK  ACIDS.    EQUILIBRIUM. 

Experiment.  Prepare  normal  and  o.i  normal  acetic  acid  solutions*  from 
the  laboratory  6-normal  solution. 

Determine  if  the  o.  i  normal  aqetic  acid  soution  has  acid  properties  by 
testing  it  with  litmus,  methyl  orange,  sodium  carbonate,  and  by  tasting  it  (Cf. 
Assignment  VI). 

Estimate  approximately  the  concentration  of  hydrogen  ion  in  the  normal 
and  o.i  normal  acetic  acid  solutions  by  testing  10  cc.  of  each  solution  with  a 
a  drop  of  methyl  violet  solution,  and  comparing  the  colors  with  those  obtained 
with  hydrochloric  acid  solution  and  water  in  the  preceding  Assignment. 
Notes  and  Questions.  The  reasoning  depends  upon  the  comparison  of  the 
lowest  concentration  of  solutions  of  the  two  acids  at  which  methyl  violet  gives  a 
color  distinctly  different  from  that  with  water.  The  acetic  acid  in  the  solu- 
tion must  be  present  either  in  the  form  of  ions,  H+  and  Ac',  or  in  the  un-ion- 
ized  form,  HAc.  From  your  estimate  of  the  concentration  of  H+  in  the  o.i 
normal  solution,  calculate  the  fraction  which  is  ionized,  and  the  fraction  which 
is  un-ionized.  The  fraction  of  the  acid  which  is  in  the  form  of  ions  is  called 
the  degree  of  ionization.  State  also  the  concentrations  of  acetate  ion,  Ac",  and 
of  the  un-ionized  acid,  HAc,  in  the  o.i  normal  acetic  acid  solution.  Is  acetic 
acid  a  weak  or  a  strong  acid? 

Questions.  The  con^ntration  of  the  ions  in  acetic  acid  solutions  have 
been  determined  more  accurately  than  is  possible  by  these  color  experiments. 

*  Save   some    0.1    normal    acetic    acid    solution    for    use    in    Assignment    XIII. 

18 


The  concentrations  of  hydrogen  ion  in  these  solutions  at  room  temperature  are 
given  in  the  following  table. 

Concentration  Concentration    Concentration  Concentration     of     Degree  of 

of    acid.  of  H+.                 of  Ac.~          un-ionized   HAc.      lonization. 

i  N  .004  N 

0.5  N  .003  N 

o.i  N  .0013  N 

Fill  in  the  remaining  columns  of  the  table.  Does  the  degree  of  ionization 
decrease  or  increase  as  the  concentration  of  the  acid  decreases? 

In  any  given  solution  containing  acetic  acid  each  of  the  substan- 
ces, H+,  Ac",  and  HAc,  has  a  definite  concentration.  The  three  substances  are 
in  equilibrium,  and  their  concentration  may  be  changed  by  altering  the  experi- 
mental conditions,  e.  g.,  by  raising  or  lowering  the  temperature,  or  by  increas- 
ing or  decreasing  the  concentration  of  one  or  more  of  the  substances  involved. 
A  careful  study  of  this  equilibrium  will  enable  the  student  to  understand  more 
readily  the  many  other  examples  of  equilibrium  that  will  be  met  with  in  this 
course. 

First  consider  the  restriction  that  these  three  substances  are  not  indepen- 
dent of  each  other, — since,  namely,  H+  and  Ac~  form  HAc  when  they  unite, 
and  HAc  forms  H+  -f-  ACT  when  it  ionizes.  It  is  important  to  realize  that 
whenever  there  is  a  disturbance  of  this  equilibrium,  either  the  ions  unite  to 
form  un-ionized  HAc,  or  HAc  breaks  up  into  the  ions.  These  ideas  are  all 
involved  in  the  statement  that 

H+  +  Ac-  =  HAc 

is  a  reversible  reaction;  and  the  problem  is  to  determine  how  the  reaction  may 
be  made  to  proceed  to  the  right  or  to  the  left. 

Questions.  What  reaction  takes  place  (a)  when  acetic  acid  gas  is  dis- 
solved in  water,  (b)  when  water  is  added  to  a  normal  solution  of  acetic  acid? 
In  the  latter  case  consider  the  degrees  of  ionization  which  you  calculated  for 
acetic  acid  at  different  concentrations. 

Experiment.  The  reaction  between  sodium  acetate  and  hydrochloric  acid. 
—Prepare  some  approximately  half  normal  hydrochloric  acid.  Measure  out 
three  15  cc.  portions  in  test  tubes,  and  add  two  or  three*  drops  of  methyl  violet 
to  each.  Measure  out  $  cc.  2  N  sodium  acetate  solution  and  add  the  solution, 
a  few  drops  at  a  time,  to  one  of  the  0.5  N  hydrochloric  acid  solutions;  shake 
the  mixture  after  each  addition,  recording  the  color  and  estimating  the  volume 
of  the  sodium  acetate  solution  added.  Repeat  the  experiment  to  check  your 
results.  What  conclusions  can  you  draw  with  regard  to  changes  in  the  hydro- 
gen ion  concentration?  (Note  that  sodium  acetate,  like  other  salts,  is  present 
in  solution  as  ions,  in  this  case  Na+  and  Ac".)  State  what  reaction  has  taken 
place  and  write  the  equation.  Could  the  result  of  this  experiment  have  been 
predicted  from  your  previous  experiments  on  acetic  acid  in  this  Assignment? 

Experiment.  The  reaction  between  sodium  acetate  and  acetic  acid. — Pre- 
dict what  will  happen  when  sodium  acetate  solution  is  added  to  acetic  acid 
solution.  Give  your  reasoning.  Test  your  answer  experimentally  by  adding 
2  N  sodium  acetate  solution  to  normal  acetic  acid  containing  a  drop  of  methyl 
violet. 

Review  the  lectures  on  equilibrium  while  answering  the  following  prob- 
lems. 

Problems.'  i.  A  vessel  is  partially  filled  with  water;  all  the  air  in  it  is 
removed,  and  the  vessel  is  closed.  What  is  the  total  pressure  in  the  vessel  at 
22°  (see  Vapor  Pressure  of  Water,  Table,  Assignment  III)?  What  will  the 
total  pressure  become  if  the  vessel  is  heated  to  30°  ?  How  has  the  weight  of  water 

19 


vapor  in  the  vessel  changed  (qualitatively)  ?  What  reaction  has  taken  place 
during  the  change  from  22°  to  30°?  If  the  vessel  is  again  brought  to  22° 
what  will  be  the  pressure?  Write  the  reaction  that  has  taken  place  during 
cooling.  Suggest  a  method  for  making  this  reaction  go  in  either  direction 
without  changing  the  temperature. 

2.  State  in  words  what  is  represented  by  each  of  the  following  equations. 
Suggest   experimental    conditions   under  which   each    reaction   can   be   made   to 
proceed    (i)    towards  the   right  and    (2)    towards   the  left.     Under  what   con- 
ditions will   equilibrium  be  established  in   each  case? 

(a)  HC1   (gas)  =  H+  +  Cl~   (in  solution) 

(b)  Zn   (solid)   +  H2O   (gas)  =  ZnO   (solid)   +  H2 

3.  Are  the  gases,  hydrogen,  oxygen,  and  nitrogen  slightly  soluble  or  very 
soluble  in  water?     Give  reasons  for  your  answers.     What  happens  when  water 
saturated    with    one    of    these    solutions    is    boiled?      How    is    the    equilibrium 
between  the  gas  phase  and  the  solution  altered  by  an  increase  of  temperature? 
What  substances  are  present  in  solutions  of  each  of  these  gases? 

ASSIGNMENT  XT. 
NEUTRALIZATION  OF  ACETIC  ACID  SOLUTION  BY  A  STRONG  BASE. 

The  purpose  of  Assignment  XI  is  to  study  the  reaction  between  solutions 
of  acetic  acid  and  of  the  strong  base,  sodium  hydroxide,  and  continue  the 
writing  of  ionic  equations. 

Question.  How  many  grams  of  acetic  acid  are  contained  in  a  liter  of 
molal  solution?  The  formula  is  HC2H3O2. 

Experiment.  Procure  two  burettes  at  the  office,  fill  one  with  your  sodium 
hydroxide  solution  of  known  concentration  (Assignment  VIII),  and  the  other 
with  the  special  acetic  acid  solution  marked  for  use  in  this  Assignment.  The 
concentration  of  the  latter  solution  is  in  the  neighborhood  of  molal ;  the  exact 
value  will  be  announced.  Titrate  the  two  solutions,  using  phenolphthalein  as 
indicator,  and  save  the  solutions  obtained  in  the  titrations.  Note  if  the  end- 
point  is  sharp,  and  if  the  results  of  two  or  more  titrations  are  concordant. 
Calculate  the  normal  concentration  of  the  acetic  acid  solution.  Questions. 
What  conclusions  do  you  draw  from  your  results?  How  many  mols  of  acetic 
acid  are  neutralized  by  i  equivalent  of 'sodium  hydroxide? 

Experiment.  Acidify  the  solution  obtained  in  one  of  the  above  titrations 
with  a  little  acetic  acid,  and  evaporate  it  to  dryness,  taking  care  not  to  heat 
the  residue  strongly  while  driving  off  the  last  portions  of  water  and  acetic 
acid.  Questions.  What  is  the  solid  residue?  From  the  results  of  the  titra- 
tions write  the  equation  (non-ionic)  for  the  neutralization  of  acetic  acid  by 
sodium  hydroxide.  Point  out  how  it  is  possible  for  this  solution  of  a  weak  acid, 
in  which  the  concentration  of  H+  is  so  small,  to  neutralize  so  much  sodium 
hydroxide.  Which  is  the  weaker  electrolyte,  acetic  acid  or  water? 

If  you  have  answered  the  questions  in  the  preceding  paragraph  correctly 
you  will  realize  that  it  is  impossible  to  represent  all  that  has  happened  in  this 
neutralization  reaction  in  a  single  equation.  However,  by  following  the  plan 
adopted  in  Assignment  IX  we  can  find  the  one  equation  that  shows  most 
satisfactorily  the  difference  between  the  initial  solutions  and  the  final  solution. 

HC2H,O2  -f   (Na+  +  OH-)  «==  H2O  +  Na+  +  C2H3O2-. 

This  equation  shows  that  NaOH  and  NaC2H2O2  are  strong  electrolytes,  that 
the  weak  acid  is  mainly  present  in  the  un-ionized  form,  and  that  the  sodium 
ion  takes  no  part  in  the  reaction.  Accordingly  the  equation 

HC2H302   +   OH-  =  H20  +  C,H302- 
20 


represents  the  main  reaction.  It  tells  nothing,  however,  about  the  "mechanism" 
of  the  reaction. 

Problems,  i.  How  many  grams  of  sodium  acetate  can  be  prepared  (a) 
from  50  cc.  normal  acetic  acid  solution,  (b)  from  100  cc.  0.5  normal  acetic  acid 
solution? 

2.     Outline   experiments   to   distinguish   between 

(a)  i.o  TV  HNO,  and  o.i  N  HNO3, 

(b)  o.oi  N  HNO3  and  i.o  N  acetic  acid. 

Optional  Experiment.  This  is  to  be  ommitted  by  the  majority  of  the 
class,  and  should  not  be  porformed  without  consulting  the  instructor.'  Repeat 
the  titration  of  the  acetic  acid  and  sodium  hydroxide  solutions,  using  methyl 
orange  as  indicator. 

ASSIGNMENT  XII. 
STRONG  AND  WEAK  BASES.     EQUILIBRIUM. 

The  purpose  of  Assignment  XII  is  to  compare  the  concentration  of 
hydroxide  ion  in  solutions  of  different  bases,  and  to  continue  the  study  of 
equilibrium. 

\ 

Experiment.  Prepare  solutions  of  sodium  hydroxide  which  are  approxi- 
mately normal,  o.i  normal,  o.oi  normal,  and  o.ooi  normal.  (State  in  your 
note-book  how  you  prepared  these  solutions.)  To  10  cc.  of  each  solution  in 
a  test  tube,  add  i  drop  of  a  solution  of  the  indicator,  trinitrobenzol.  Record 
the  color  obtained  in  each  case,  and  note  carefully  the  most  dilute  solution  that 
gives  a  color  with  the  indicator. 

Repeat  the  experiment  with  potassium  hydroxide  solution,  and  compare 
the  colors  obtained  at  each  concentration.  If  sodium  hydroxide  in  solution  is 
completely  ionized,  what  conclusion  can  you  draw  with  respect  to  potassium 
hydroxide?  What  concentrations  of  hydroxide  ion  can  be  measured  by  means 
of  this  indicator,  trinitrobenzol? 

Try  similar  experiments  with  ammonium  hydroxide  solutions  of  various 
concentrations,  say  normal  and  o.i  normal*.  Point  out  the  sodium  hydroxide 
solution  and  the  ammonium  hydroxide  solution  which  give  approximately  the 
same  faint  color  with  the  indicator,  and  estimate  the  concentration  of  hydrox- 
ide ion  in  this  ammonium  hydroxide  solution.  What  is  the  concentration  of 
ammonium  ion  NH4+  in  this  solution?  What  is  the  concentration  of  the  un- 
ionized ammonium  hydroxide,  NH4OH?  What  is  the  "degree  of  ionization"? 
Is  ammonium  hydroxide  a  strong  or  a  weak  base? 

In  every  solution  of  ammonium  hydroxide,  just  as  in  the  case  of  acetic 
acid,  there  is  a  definite  equilibrium  between  the  un-ionized  substance  and  the 
two  ions.  This  equilibrium  is  represented  by  the  equation 

NH4OH  =  NH4+  +  OH-. 

Taking  into  account  the  results  obtained  in  these  experiments,  predict 
what  will  happen  when  solutions  of  sodium  hydroxide  and  ammonium  chloride 
are  mixed.  Test  your  answer  by  means  of  an  experiment  with  these  solutions 
and  the  indicator,  trinitrobenzol.  Write  an  equation  for  the  reaction  that 
takes  place. 

How  will  the  equilibrium  in  an  ammonium  hydroxide  solution  be  disturbed 
when  ammonium  chloride  is  added?  What  reaction  will  take  place?  Test 
your  answer  by  an  experiment  with  i  normal  NH4OH,  ammonium  chloride 
solution  and  trinitrobenzol. 

Problems.  i.  An  ammonium  hydroxide  solution  has  a  characteristic 
odor;  this  is  due  to  the  NH3  gas  given  off  by  the  solution.  Which  solutions 

*Save  some  0.01  normal  NaOH  and  0.1  normal  NH  OH  solutions  for  use  in  Assignment  XIII. 

4 

21 


smell  more  strongly  of  ammonia,  (a)  a  dilute  solution  or  a  concentrated  solu- 
tion? (b)  a  cold  solution  or  a  hot  solution?  Write  an  equation  to  show  the 
formation  of  NH4OH  from  NH3  and  water.  Is  this  a  reversible  reaction?  Is 
there  an  equilibrium  between  gaseous  NH3  and  the  solution?  Give  reasons  for 
your  answers.  In  this  discussion  point  out  what  happens  (a)  when  NH3  gas 
is  passed  into  pure  water,  (b)  when  the  resulting  solution  is  boiled  in  an  open 
vessel,  (c)  when  the  gas  space  above  a  solution  is  evacuated. 

2.  How  will  the  equilibrium  NH4OH  :  :  NH4+  -f  OH-  be  disturbed 
when  sulfuric  acid  is  added  to  an  ammonium  hydroxide  solution?  State 
how  you  would  prepare  ammonium  sulfate  from  ammonium  hydroxide  and 
sulfuric  acid.  Write  an  equation  to  represent  the  main  reaction  in  aqueous 
solution. 

ASSIGNMENT  XIII. 
HYDROLYSIS. 

In  earlier  Assignments  neutralization  reactions  have  been  shown  to  depend 
upon  the  fact  that  the  reaction  H+  -(-  OH~  =  H2O  takes  place  practically 
completely.  The  concentration  of  H+  and  OH"  in  pure  water  is,  however,  not 
zero;  each  has  a  definite,  although  very  small,  value,  namely  icr7  N  (i/io7). 
Water  is  accordingly  a  very  weak  electrolyte ;  in  other  words,  the  reaction 
between  H+  and  OH~  does  not  go  quite  to  completion,  and  there  is  in  all 
aqueous  solutions  an  equilibrium 

H,O  =  H+  -f-  OH-. 

In  Assignment  XIII  we  shall  study  some  experiments  closely  connected  with 
this  fact. 

A  solution  of  sodium  acetate  may  be  prepared  in  two  ways  :  ( i )  by  mixing 
exactly  equivalent  amounts  of  acetic  acid  and  sodium  hydroxide  (reaction, 
HAc  +  Na+  +  OH-  =  Na+  -f  Ac-  +  H2O,  or  more  simply,  HAc  +  OH- 
=  Ac"  -j-  H2O),  and  (2)  by  dissolving  solid  sodium  acetate  in  water.  Both 
solutions  must  have  exactly  the  same  properties. 

Prepare  some  O.OT  N  NaOH,  o.oi  N  HC1,  o.i  N  acetic  acid,  and  o.i  N 
ammonium  hydroxide  for  use  in  the  following  experiments. 

Experiment.  To  10  cc.  4  N  sodium  acetate  solution  add  one  drop  litmus 
solution.  To  another  10  cc.  portion  of  the  solution  add  one  drop  phenol- 
phthalein  solution.  Repeat  with  10  cc.  portions  of  distilled  water.  What 
results  did  you  obtain  in  similar  experiments  with  NaCl  solution  in  Assign- 
ment VII?  (Repeat  these  experiments  if  necessary.) 

Questions.  In  what  respect  does  the  sodium  acetate  solution  differ  in 
properties  from  water  and  from  a  sodium  chloride  solution?  Is  the  reaction 
between  equivalent  amounts  of  acetic  acid  and  sodium  hydroxide  complete? 
What  substances  are  present  in  the  resulting  solution?  State  what  happens 
when  solid  sodium  acetate  is  dissolved  in  water.  The  reaction  that  takes  place 
between  the  acetate  ion  and  water  is  an  example  of  hydrolysis.  Write  this 
reaction  as  well  as  you  can  in  a  single  equation,  and  compare  the  equation 
with  that  given  above  for  the  neutralization  of  acetic  acid  by  sodium  hydrox- 
ide. Which  of  the  following  substances  in  the  4  N  sodium  acetate  solution 
have  (i)  large  concentrations,  (2)  small  concentrations,  and  (3)  extremely 
small  concentrations:  H2O,  Na+,  Ac",  HAc,  OH~,  H+.  Test  your  answer  by 
the  following  experiments : 

Experiment.  To  10  cc.  4  N  sodium  acetate  solution  containing  i  drop 
phenolphthalein  add  o.i  N  acetic  acid,  drop  by  drop.  State  in  words  what 
happens  in  this  reaction.  Write  the  equation  for  the  main  reaction. 

Experiment.  To  10  cc.  distilled  water  containing  i  drop  phenolphthalein 
add  o.oi  N  NaOH  solution  drop  by  drop  until  (after  shaking)  the  color  is  the 

22 


same  as  in  the  4  N  NaAc  solution.  Estimate  the  volume  of  one  drop  by 
counting  the  number  of  drops  in,  say  5cc.,  and  calculate  the  concentration  of 
hydroxide  ion  in  4  N  sodium  acetate.  What  is  the  concentration  of  un-ionized 
acetic  acid  in  the  same  solution?  What  is  the  concentration  of  Na+  and  of 
Ac"?  (Assume  complete  ionization  of  sodium  acetate  and  sodium  hydroxide.) 

Questions.  Do  these  results  furnish  any  evidence  that  the  reaction  H+  -)- 
OH"  =  H2O  is  reversible?  How  is  this  equilibrium  disturbed  by  the  addition 
of  sodium  hydroxide?  Is  the  concentration  of  hydrogen  ion  in  the  4  N 
sodium  acetate  solution  greater  or  less  than  in  pure  water? 

Experiment.  Repeat  the  above  experiments  with  10  cc.  4  TV  ammonium 
chloride.  State  whether  the  solution  has  an  acid  or  basic  reaction. 

To  10  cc.  of  4  N  ammonium  chloride  containing  a  drop  of  litmus  solution 
add  gradually  o.i  N  NH4OH  until  the  solution  is  neutral.  What  reaction  has 
taken  place?  Discuss  the  hydrolysis  of  ammonium  chloride  and  the  neutraliz- 
ation of  the  weak  base  ammonium  hydroxide  by  an  equivalent  amount  of 
hydrochloric  acid.  Write  a  single  equation  to  represent  the  main  reaction  in 
each  case.  Devise  and  try  an  experiment  to  determine  the  hydrogen  ion 
concentration  in  2  N  ammonium  chloride. 

Questions.  How  do  the  concentrations  of  H+  and  OH~  in  pure  water  and 
in  solutions  of  NaCl,  NaNO3,  K2SO4,  compare  with  each  other?  What  con- 
clusions can  you  draw  with  regard  to  (a)  the  hydrolysis  of  the  salt  of  a  strong 
acid  and  a  strong  base,  (b)  the  hydrolysis  of  the  salt  of  a  weak  acid  and  a 
strong  base,  as  potassium  acetate,  (c)  the  hydrolysis  of  the  salt  of  a  strong 
acid  and  a  weak  base,  as  ammonium  nitrate? 

State  whether  or  not  you  would  expect  a  salt  of  a  weak  acid  and  a  weak 
base  (as  ammonium  nitrate)  to  be  hydrolyzed.  Ammonium  acetate,  like  other 
salts,  is  completely  ionized  in  dilute  solution.  When  solid  ammonium  acetate 
is  dissolved,  what  reaction  will  take  place  between  ( i )  ammonium  ion  and 
water,  (2)  between  acetate  ion  and  water?  Write  a  single  equation  to  repre- 
sent the  hydrolysis.  Will  the  neutralization  reaction  between  equivalent 
amounts  of  acetic  acid  and  ammonium  hydroxide  take  place  completely?  Will 
it  take  place  more  or  less  completely  than  the  neutralization  of  acetic  acid  by 
sodium  hydroxide? 

Experiment.  Test  a  normal  solution  of  ammonium  acetate  with  indicat- 
ors. Explain  your  results,  taking  into  account  the  fact  that  acetic  acid  and 
ammonium  hydroxide  are  ionized  to  about  the  same  extent. 

Problems,  i.  In  the  table  given  below  fill  in  the  formulas  of  the  com- 
pounds which  the  given  positive  and  negative  ions  form  with  each  other. 

Now  mark  with  a  star  (*)  those  compounds  which  are  only  slightly  ion- 
ized in  solution,  and  mark  with  an  "h"  those  compounds  which  are  hydrolyzed 
in  solution. 

Study  this  table  carefully  and  point  out  parallelisms  between  the  strengths 
of  acids  and  bases  and  the  hydrolysis  of  the  corresponding  ions. 

H+  Na+  K+  NH4+ 

OH- 

ci- 

NO.,- 

so4- 

Ac- 

2.  Summarize  the  behavior  of  solutions  towards  the  five  indicators  you 
have  used  in  Assignments  VI  to  XIII  by  completing  the  following  table. 

23 


Concentration.  Methyl  violet.  Methyl  orange. 

H+  i.o       normal  yellow  red 

H+  o.i 
H+  o.oi 

H+  o.ooi  violet 

H+  slightly  greater  than  in  red 

water  yellow 

H+  and  OH~  equal  in 

water 
OH"  slightly   greater  than 

in  water 

OH~  o.oi  normal 
OH-  o.i 
OH-  i.o 

State  how  you  would  estimate   roughly  the  H+   cr   OH"   concentration   in   any 
colorless  unknown  solution.     Suggest  a  definition  for  an  indicator. 

ASSIGNMENT  XIV. 
CARBON  DIOXIDE,  CARBONATES,  BICARBONATES,  CARBONIC  ACID. 

In  Assignments  VI  and  X  solutions  of  acids  were  found  to  react  with 
sodium  carbonate  to  give  a  colorless,  odorless  gas.  This  gas  is  carbon  dioxide, 
CO2.  In  Assignment  XIV  we  shall  study  this  reaction  more  carefully,  .and 
shall  investigate  the  behavior  of  an  aqueous  solution  of  carbon  dioxide. 

Experiment.  Add  a  small  quantity  of  powdered  sodium  carbonate, 
Na2CO3,  to  10  cc.  water,  shake  the  mixture  several  times,  and  filter  if  there  is 
a  residue.  Add  some  of  the  clear  solution,  a  little  at  a  time,  to  a  solution  of 
(a)  hydrochloric  acid,  (b)  acetic  acid.  Is  Na2CO3  a  readily  soluble  or  a 
difficultly  soluble  substance?  Like  other  salts,  it  is  highly  ionized  in  aqueous 
solution.  Complete  the  equations 

H+  +  CO,-  =  CO,  -f 
HAc  +  CO3~  =  CO2  + 

A  Test  for  Carbon  Dioxide.  Moisten  the  end  of  a  glass  tube  with  Ba( OH) 2 
solution  and  hold  it  over  a  solution  from  which  CO2  bubbles  are  issuing 
(or  pass  CO2  gas  into  i  or  2  cc.  Ba(OH)2  solution).  The  white  solid  is 
barium  carbonate,  BaCO3,  and  its  formation  from  Ba(OH)2  solution  is  a  test 
for  carbon  dioxide.  Prove  that  the  breath  contains  CO2.  Is  BaCO,  readily 
soluble  or  difficultly  soluble  in  water?  Complete  the  equation 

CO2  +  Ba++  +  OH-  =  BaCO3  + 

Note  that  Ba(OH)2  is  a  strong  base.     Outline  an   experiment   to  prove   this. 

Experiment.       Repeat    the    first    experiment,    using    sodium    bicarbonate, 

NaHCO,,   instead   of   Na2CO3.      Report   whether   it  is    readily   soluble   or   diffi-. 

cultly   soluble   in   water.      Prove   that   the   gas   evolved   is   CCX.      Complete   the 

equati°ns  H+  +  HC03-  =  CO.  + 

HAc  +  HCO3-  =  CO2  + 

and    compare    them    with   the   equations    for   the    corresponding   reactions    with 
ca'rbonate  ion. 

Experiment.  Limestone  is  calcium  carbonate,  CaCO3.  Is  it  readily  or 
difficultly  soluble  in  water?  Try  an  experiment,  if  necessary.  Test  the  action 
of  hydrochloric  and  acetic  acid  on  small  amounts  of  solid  calcium  carbonate, 
and  complete  the  equations 

H+  -f-  CaCO3   (solid)  = 
HAc  +  CaCO3   (solid)  = 

24 


State   what  you   think  will  happen   when  solid   BaCO3   is   treated  with  hydro- 
chloric acid.   Try  the   experiment  and  write  the  equation   for  the  reaction. 

Set  up  a  carbon  dioxide  generator  similar  to  that  used  in  preparing  hydro- 
gen  (Assignment  V)  ;  sign  a  white  slip  for  the  thistle  tube.     Place  in  the  flask 
a  few  small  lumps    (about  5  grams)   of  limestone,  'cover  with  water,  and  add 
a  little  hydrochloric  acid  through  the  thistle  tube. 
Show  by  experiments : 

Whether  CO2  is  denser  or  lighter  than  air. 
Whether  it  is  inflammable  or  supports  combustion. 
Whether  or  not  it  dissolves  in  water.     Give  details  of  the  last  experi- 
ment. 

The  solution  of  carbon  dioxide  in  water  contains  carbonic  acid,  H2CO3. 
Prove  by  experiment  that 

CO2   (gas)    +  H2O  =  H2CO3   (in  solution) 

is  a  reversible  reaction,  and  that  therefore  is  an  equilibrium.     At  room  tempera- 
ture  a   solution   in   equilibrium   with   CO2   gas   at    i    atmosphere   pressure   con- 
tains about  0.04  mol  H2CO3  in  i  liter.     How  is  this  equilibrium  altered  by  an  ' 
increase  of  temperature? 

Properties  of  a  Solution  of  Carbonic  Acid.  Prepare  a  small  quantity  of 
a  nearly  saturated  solution  of  CO2  and  test  it  with  indicators  (see  Problem 
2,  Assignment  XIII).  Is  H2CO3  a  weak  or  a  strong  acid?  Decide  if  possible 
whether  it  is  weaker  or  stronger  than  acetic  acid.  The  equilibrium  in  the 
solution  depends  on  the  reversible  reaction 

H2CO3  =  H+  -f  HCO3-. 

Explain  what  happens  in  the  following  experiments : 

To  10  cc.  water  add  i  drop  litmus  solution,—] ust  enough  to  give  a  faint 
bluish  color.  Pass  in  CO2  gas  until  the  color  changes.  Heat  the  solution  to 
boiling  and  cool  it  again. 

To  10  cc.  water  containing  i  drop  phenolphthalein  add  a  drop  of  dilute 
NaOH  solution, — just  enough  to  give  a  faint  pink  color;  pass  in  CO,  gas. 

Review  the  above  experiments  on  the  action  of  acids  on  NaHCO3  solution, 
Na2COo  solution,  and  solid  CaCO3,  and  state  what  happened  in  the  solution 
before  CO2  gas  was  evolved. 

The  Neutralization  of  Carbonic  Acid  in  Steps.  Which  of  the  following 
equations  represents  the  main  reaction  when  solutions  of  equimolal  quantities 
of  H2CO,  and  NaOH  are  mixed? 

H2CO3  -f  OH-  =  HCO3-  +  H,O 
H+  -f  OH-  =  H2O 

Experiment.  The  reaction  between  bicarbonate  ion  and  hydroxide  ion. — 
To  a  solution  of  sodium  bicarbonate  containing  trinitrobenzol,  add  slowly  from 
a  graduated  cylinder  0.5  normal  NaOH  solution.  Perform  a  blank  experiment 
(for  comparison)  by  adding  the  NaOH  solution  to  an  equal  volume  of  water 
containing  the  same  indicator.  Have  you  any  evidence  that  the  reaction 

HCO3-  +   OH-  =  H2O  -f-  CO3~ 

has  taken  place?  Is  it  correct  to  say  that  bicarbonate  ion  is  an  acid?  Try  to 
devise  an  experiment  to  test  if  the  reaction  just  written  is  reversible. 

Write  the  equation  for  the  main  reaction  between  H2CO3  and  excess 
sodium  hydroxide  solution. 

Experiment.  The  hydrolysis  of  carbonate  ion. — Test  10  cc.  portions  of 
normal  Na2CO3  solution  with  indicators.  To  the  portion  containing  phenol- 
phthalein add  o.i  normal  hydrochloric  acid  very  slowly  from  a  graduate. 
Compare  your  results  with  those  obtained  with  sodium  acetate  solutions  in 
Assignment  XIII,  and  state  which  solution  has  the  greater  concentration  of 
OH~.  The  reaction  between  carbonate  ion  and  water  is 

25 


CO3~  +  H2O  =  HCO3-  -f  OH-, 

and  this  experiment  furnishes  a  proof  that  HCO3~  is  a  very  weak  acid.  Which 
is  the  weaker  acid,  acetic  acid  or  bicarbonate  ion  ?  Give  your  reasoning. 

It  is  evident  from  the  preceding  experiments  that  the  weak  dibasic  acid 
H,CO3  ionizes  in  two  stages,  namely, 

H,CO,  =  H+  +  HCO,-,  and 
HCCV  =  H+  +  CO,-. 

The  second  acid,  HCO3~,  is  a  much  weaker  acid  than  the  first,  H2CO:;.  Sug- 
gest experiments  or  give  reasons  to  prove  this  fact.  There  are  two  series  of 
salts :  the  "normal  salts"  as  Na2CO3  and  CaCO3  and  the  "acid  salts"  as 
NaHCO:i.  The  chief  negative  ions  present  in  solutions  of  these  two  types  of 
salts  are  CO3~~  and  HCO3~  respectively.  Sodium  bicarbonate  is  also  named 
sodium  acid  carbonate  or  sodium  hydrogen  carbonate ;  and  the  bicarbonate  ion 
is  named  similarly  the  acid  carbonate  ion  or  the  hydrogen  carbonate  ion. 

The  behavior  of  other  weak  polybasic  acids  is  similar  to  that  of  carbonic 
acid.  Thus  the  weak  dibasic  acid,  hydrogen  sulfide  in  solution,  H2S,  ionizes 
in  two  stages,  and  there  are  two  series  of  salts.  The  second  acid.  HS~,  is 
weaker  than  the  first,  H,S. 

Sulfuric  acid,  H2SO4,  is  an  example  of  a  strong  polybasic  acid.  Review 
your  experiments  with  H2SO4  in  Assignments  VIII,  and  IX,  and  decide 
whether  hydrogen  sulfate  ion,  HSO4~,  is  a  strong  or  a  weak  acid.  What  ions 
are  present  in  large  quantity  in  a  dilute  solution  of  potassium  acid  sulfate, 
KHSO4?  Try  experiments  if  necessary. 

Problems,  i  What  general  statement  can  you  make  about  the  action  of 
an  acid  on  a  solution  of  a  salt  of  a  weaker  acid?  Give  examples  and  write 
equations. 

2.  A  solution  of  sodium  bicarbonate  is  prepared  by  dissolving  8.4  grams 
of  this  salt  in  enough  water  to  give  i  liter  of  solution,  (a)  Calculate  approxi- 
mately the  concentration  of  HCO3~  in  this  solution.  (b)  What  weight  of 
sodium  carbonate  can  be  prepared  from  this  solution?  (c)  How  many  cc. 
of  normal  NaOH  solution  are  necessary  in  (b)  ?  What  volume  of  CO2  under 
standard  conditions  can  be  prepared  from  this  quantity  of  NaHCO3? 

Optional  Experiment.     Test  a  solution  of  NaHCO3  with  indicators. 

Considering  the  facts  that  H2CO3  is  a  stronger  acid  than  HCO;r,  predict 
what  will  happen  when  CO2  is  passed  into  a  solution  of  NaXCX.  Test  your 
answer  experimentally.  Write  the  equation  for  the  reaction. 

Predict  what  will  happen  when  a  solutoin  of  NaHCO3  is  boiled.  How 
would  the  concentration  of  OH~  change  during  the  boiling?  Test  your  answer 
by  experiments.  Write  the  equation  or  equations.  Is  the  reaction  considered 
in  the  preceding  paragraph  reversible? 

Note.  When  you  have  completed  Assignment  XIV,  clean  your  two  sample 
bottles,  label  them  with  your  desk  number  and  name,  mark  them  No.  i  and 
No.  2  respectively,  and  deposit  them  at  the  office.  Your  two  unknowns  (for 
analysis)  will  be  ready  for  you  at  the  next  laboratory  period.  See  Assign- 
ment XV. 

ASSIGNMENT  XV. 
REVIEW.     ANALYSIS  OF   SOLUTIONS   FOR   IONS  ALREADY   STUDIED. 

Give  formulas  and  names  of  four  or  five  of  the  common  sodium,  potas- 
sium and  ammonium  salts.  Are  these  salts  readily  soluble  or  difficultly 
soluble  in  water?  Are  all  of  them  highly  ionized  in  dilute  solution? 

What  is  the  color  of  each  of  the  ions,  Na+,  K+,  NH4+?  Give  other  prop- 
erties of  these  ions,  and  specify  a  characteristic  property  of  each  that  may  be 
used  as  a  test.  Devise  methods  of  testing  for  each  of  these  ions  in  the  pres- 
ence of  the  other  two.  Try  your  methods  with  known  solutions :  for  example, 
divide  a  solution  contining  Na+  and  NH4+  into  two  parts,  to  one  part  add  a 

26 


small  quantity  of  a  potassium  salt  solution,  and  be  sure  that  you  can  distin- 
guish between  these  two  solutions.  Is  it  Na+  or  NH4+  that  causes  difficulty 
in  testing  for  K+? 

Suggest  a  method  of  testing  a  solution  for  the  presence  of  a  carbonate 
based  on  the  evolution  and  detection  of  CO2.  Try  your  method. 

State  how  you  would  test  a  solution  for  the  presence  of  Cl~,  for  the 
presence  of  SO4".  Write  the  ionic  equations.  Apply  these  tests  to  a  solu- 
tion of  a  carbonate  free  from  Cl~  or  SO4~  in  order  to  determine  if  carbonate 
interferes  with  your  tests.  If  it  does,  suggest  and  try  a  method  of  preventing 
this  interference. 

Experiment.  Heat  to  boiling  a  normal  solution  of  (NH4)2SO4  in  a  small 
beaker,  and  note  if  there  is  an  odor.  Now  add  i  cc.  6  normal  NaOH,  and 
again  heat  to  boiling.  While  you  are  heating  the  mixture  note  the  odor  and 
test  the  vapor  with  moist  litmus.  Questions.  What  reaction  takes  place  (a) 
when  cold  solutions  of  (NH4)2SO4  and  NaOH  are  mixed?  (b)  when  the 
solution  is  boiled?  What  volume  of  6  normal  NaOH  would  be  needed  to 
react  completely  with  5  cc.  normal  (NH4)2SO4?  If  a  mixture  consisting  of 
5  cc.  normal  (NH4)2SO4  and  3  cc.  6  normal  NaOH  were  evaporated  to  dry- 
ness,  what  substances  would  be  present  in  the  solid  residue? 

Experiment.  Mix  5  cc.  normal  NaCl  and  3  cc.  6  normal  H2SO4  in  your 
casserole,  evaporate  the  mixture  on  the  porch*  until  heavy  white  fumes  of 
sulfuric  acid  are  given  off.  When  the  residue  is  cold,  pour  it  into  cold  water, 
and  test  the  solution  for  chloride.  Questions.  Did  any  reaction  take  place 
(a)  when  the  two  solutions  were  mixed,  (b)  when  the  mixture  was  evapor- 
ated? Which  is  the  more  volatile,  HC1  or  H2SO4? 

Test  the  volatility  of  nitric  acid  and  acetic  acid  by  evaporating  solutions  of 
these  acids  in  a  porcelain  dish  on  the  porch. 

Salts  of  sodium  and  potassium,  and  of  metals  in  general,  are  not  easily 
volatilized.  Test  the  volatility  of  ammonium  salts,  as  NH4C1  and  (NH4)2SO4 
by  heating  small  quantities  of  the  solids  in  a  porcelain  dish  on  the  porch. 

Analysis  Nos.  i  and  2.     Analyze  the  two  "unknowns"  for 
H+,       Na%       K+,  NH4+, 
OH-,    C1-,  S04-    C03- 

In  these  and  all  later  analyses  estimate  roughly  the  concentration  of  H+  or  OH~ 
in  the  original  solution  ^see  Problem  2,  Assignment  XIII). 

(Optional  work.     Distinguish   carbonates   and  bicarbonates.) 
While   making   these   analyses    record   in   your   note-book   all   your  experi- 
ments   and    observations,    and    state    clearly    the    conclusions    that    you    draw. 
Submit   this   written    report   of   your   analyses    to   your   instructor   before   pro- 
ceeding with   the  next  Assignment. 

Problems,     i.     State  how  you  would  convert: 

(a)  Solid  sodium  carbonate  into  solid  sodium  nitrate. 

(b)  Solid  sodium  nitrate  into  solid  sodium  sulfate. 

(c)  Solid  ammonium  chloride  into  solid  sodium  chloride. 

2.  How  would  you  recover  the  potassium  chloride  from  a  mixture  of 
potassium  chloride  and  ammonium  chloride. 

ASSIGNMENT  XVI. 
CHEMISTRY  OF  CALCIUM. 

In  Assignment  XVI  we  shall  study  the  chemistry  of  metallic  calcium  and 
of  calcium  ion.  We  shall  find  that  certain  compounds  of  this  element  differ 
from  the  corresponding  compounds  of  sodium,  potassium,  and  ammonium  in 

*  All     operations    which     give     rise     to    noxious     fumes     must     be    performed     on     the 
porch,    not    in    the    laboratory. 

27 


that  they  dissolve  to  a  much  smaller  extent  in  water;  and  we  shall  study  the 
equilibrium  between  these  solids  and  their  "saturated"  solutions.  Review  the 
lectures  on  calcium.) 

The  solubility  of  a  substance  in  water  is  the  concentration  of  its  satur- 
ated solution.  Question.  If  the  solubiltiy  of  NaCl  in  water  at  room  tempera- 
ture is  5.4  mols  per  liter,  how  many  grams  of  NaCl  are  contained  in  100  cc. 
of  saturated  solution?  How  could  you  prove  that  NaCl  (solid)  =  NaCl  (in 
solution)  is  a  reversible  reaction? 

Note.     In  future  write  equations  for  all  reactions. 

Obtain  from  the  office  a  piece  of  metallic  calcium.  Describe  its  proper- 
ties as  far  as  you  can  observe  them  by  physical  examination.  In  what 
respects  does  calcium  show  the  physical  properties  of  a  metal? 

Experiment.  Drop  the  calcium  into  20  cc.  of  water.  Stir  the  mixture  or 
warm  gently  until  the  metal  has  dissolved.  Test  the  solution  for  OH".  What 
reaction  occurs?  The  white  solid  formed,  calcium  hydroxide,  is  a  strong  base 
but  only  slightly  soluble.  Question.  Name  other  metals  that  react  readily 
with  water?  What  reactions  occur? 

Write  the  ionic  reaction  for  the  dissolving  of  solid  calcium  hydroxide  to 
form  a  saturated  solution,  and  indicate  the  equilibrium  involved.  How  can 
the  equilibrium  be  shifted  so  that  more  Ca(OH)2  will  dissolve?  So  that  less 
will  dissolve? 

Experiment.  Divide  the  solution  of  Ca(OH)2  and  suspended  solid  into 
three  parts.  Set  aside  one  of  these  (corked)  for  future  use.  To  one  part 
add  10  cc.  of  i  N  NH4C1.  Heat  the  solution  and  test  the  odor.  To  another 
part  add  3  cc.  of  6  N  HC1.  What  reactions  have  taken  place  in  these  two 
cases?  Note  the  relation  of  these  reactions  to  the  equilibrium  discussed  in 
the  preceding  paragraph. 

To  5  cc.  i  N  CaCU  add  Na2CO3  solution  in  excess,  heat  the  mixture  to 
boiling,  filter.  (The  instructor  will  explain  the  details  of  the  process  of  filtra- 
tion.) The  white  precipitate  is  calcium  carbonate.  Write  the  simplest  ionic 
equation  for  its  formation.  Prove  that  the  fiiltrate  contains  chloride  ion. 

The  fact  that  the  amount  of  calcium  ion  in  the  filtrate  must  be  very  small 
can  be  deduced  from  the  small  solubility  of  CaCO3  in  water,  0.00013  mols  per 
liter. 

In  any  solution  saturated  with  solid  calcium  carbonate  there  is  an  equili- 
brium between  the  solid  substance  and  its  ions,  which  may  be  expressed  by 
the  equation 

CaCO3  (solid)  =  Ca++  -f  CO3". 

When  will  a  precipitate  of  calcium  carbonate  be  formed?  How  is  this  equi- 
librium disturbed  by  the  addition  of  a  strong  acid?  Review  or  repeat  your 
experiments  (Assignment  XIV)  on  the  reaction  of  hydrochloric  and  acetic 
acids  on  CaCO3.  Which  is  the  stronger  acid,  HAc  or  H2CO3? 

Experiment.  Pass  CO2  gas  into  a  solution  of  CaCL.  What  conclusion 
do  you  draw  with  regard  to  the  concentration  of  CO3~  in  a  solution  of 
carbonic  acid? 

To  a  solution  of  CaCl2  add  NaOH  solution  slowly  a  few  drops  at  a  time. 
Repeat  the  experiment  with  NH4OH  solution.  What  conclusions  do  you 
draw  with  regard  to  the  relative  concentrations  of  OH"  in  NaOH  and  NH4OH 
solution? 

Now  pass  CO2  gas  into  the  CaCl2  solution  containing  NH4OH,  and  warm 
the  mixture  gently. 

Filter  the  third  portion  of  the  mixture  obtained  by  the  action  of  calcium 
on  water,  and  pass  a  few  bubbles  of  CO2  gas  into  the  filtrate.  Under  what 
conditions  will  CO2  gas  form  a  precipitate  of  CaCO3  in  a  solution  containing 
Ca++?  Which  is  the  less  soluble  substance,  CaCO3  or  Ca(OH)2? 

28 


Test  the  action  of  excess  CO2  on  the  same  mixture  by  passing  CO2  gas 
into  a  small  quantity  of  the  mixture  for  some  time.  The  reaction  involves 
the  formation  of  bicarbonate  ion,  HCO3",  H2CO3  +  CaCO3  (solid)  =  Ca++  -f- 
2HCO3~.  There  is  an  equilibrium.  Point  out  how  the  reverse  reaction  can 
be  made  to  take  place.  (Optional  Question.  Point  out  how  these  results  are 
related  to  the  results  of  the  optional  experiments  at  the  end  of  Assignment 
XIV.) 

Experiment.  Determine  whether  calcium  salts  give  a  characteristic  flame 
test.  Can  you  distinguish  the  flame  with  a  calcium  salt  from  that  with  sodium 
or  potassium  salts? 

Experiment.  To  10  cc.  normal  CaCl2  solution  add  2  cc.  6  normal 
HoSO4.  If  no  precipitate  appears  at  once,  heat  the  solution  gently,  and  let  it 
stand.  Filter.  Test  a  portion  of  the  filtrate  for  Ca++  by  adding  NH4OH 
until  the  solution  is  no  longer  acid,  and  then  (NH4)2CO3  solution,  and  warm 
the  mixture.  Test  another  portion  for  SO4  by  adding  BaCL  solution.  What 
conclusion  do  you  draw  with  regard  to  the  solubility  of  CaSO4?  Which  is 
the  less  soluble,  (i)  CaSO4  or  CaCO3,  (2)  CaSO4  or  BaSO4? 

Predict  what  will  take  place  when  solid  CaSO4  is  heated  with  excess 
NaXO3  solution?  Experiment.  Test  your  prediction  by  boiling  some  solid 
CaSO/with  normal  Na,CO3  solution/  Test  the  filtrate  for  SO4".  Wash 
the  precipitate  with  water,  and  test  it  for  carbonate. 

Test  whether  the  reverse  reaction  will  take  place  by  heating  solid  CaCO3 
with  sodium  sulfate  solution,  filtering  and  testing  the  precipitate  and  filtrate. 

Problems.  i.  Arrange  the  compounds  of  calcium  according  to  their 
solubilities  in  water,  distinguishing  readily  soluble,  moderately  soluble,  and 
difficultly  soluble  substances.  Point  out  the  compounds  which  are  much  more 
soluble  in  dilute  hydrochloric  or  nitric  acids  than  in  water. 

2.  Can  the  following  substances  be  present  at  high  concentrations  in  the 
same  solution?  If  not,  what  is  formed? 

(a)  H+  and  NO,-  (e)  H+  and  HCO3- 

(b)  H+  and  OH-  (f)  Ca++  and  CO3~ 

(c)  H+  and  S04-  (g)  Ag+  and  Cl~ 

(d)  H+  and  CO3~  (h)  H2CO3  and  CO3" 

ASSIGNMENT    XVII. 

ACTION   OF  ACTDS  ON   METALS.      SOLUBLE   SALTS  OF   ZINC, 
COPPER  AND  SILVER. 

Describe  the  physical  properties  of  zinc,  copper  and  silver,  noting  any 
characteristic  property  of  each  metal. 

Zinc  and  Acids.  Review  your  experiments  on  the  action  of  acids  on 
zinc.  Write  the  ionic  equations  for  the  reaction  between  zinc  and  solutions 
of  hydrochloric  and  sulfuric  acids.  What  is  the  color  of  Zn++?  Is  it  correct 
to  say  that  an  atom  of  the  metal  has  combined  with  2d~  or  with  SO4"?  How 
can  solid  zinc  chloride  and  solid  zinc  sulfate  be  prepared  from  these  solu- 
tions ? 

From  an  examination  of  the  ionic  equation  which  you  have  just  written 
state  how  you  would  expect  the  speed  of  the  reaction  to  alter  with  the  concen- 
tration of  H+.  Experiment.  Test  your  answer  by  treating  small  pieces  of 
zinc  with  (a)  a  small  volume  of  6  normal  HC1,  (b)  the  same  amount  of  HC1 
in  a  large  volume  of  water,  and.  (c)  6  normal  acetic  acid  solution.  In  the 
same  experiments  determine  how  the  speed  of  the  reaction  is  altered  by  an 
increase  of  temperature. 

When  Zn  reacts  with  6  normal  HNO3  the  principal  reaction,  written  in 
its  simplest  form,  is 

29 


3Zn(  solid)   +  2NO3-  -f  8H+  ==  3Zn++  -f  2NO(gas)   -f  4H2O 

NO  gas  is  colorless.  What  is  the  reaction  between  NO  and  oxygen  of  the  air? 
Repeat  the  experiment  with  Zn  and  HNO3  solution,  if  necessary.  Place  the 
metal  and  acid  in  a  small  beaker  and  cover  it  with  a  watch  glass. 

Copper  and  Acids.  Experiment.  Try  the  action  of  6  normal  HC1  and 
of  6  normal  H2SO4  on  metallic  copper.  After  a  rew  minutes  remove  the 
pieces  of  metal,  wash  them  with  water,  and  note  if  the  appearance  has  changed. 
Then  test  the  action  of  fresh  solutions  of  the  acids  on  the  clean  metal.  Is 
there  any  evidence  of  chemical  action? 

Try  the  action  of  6  normal  HNO3  on  metallic  copper.  Save  the  solution. 
What  is  the  color  of  the  copper  ion  (cupric  ion),  Cu++?  The  reaction  is 
similar  to  that  between  Zn  and  nitric  acid  solution.  Write  the  equation. 

Silver  and  Acids.  Experiment.  Try  the  action  of  6  normal  HC1,  6 
normal  PLSO.,,  and  6  normal  HNO3  on  silver.  Compare  with  the  correspond- 
ing actions  in  the  case  of  zinc  and  copper.  Save  any  solution  you  obtain. 
What  is  the  color  of  silver  ion,  Ag+?  Write  the  equation  for  the  reaction 
with  nitric  acid  solution,  noting  that  the  change  from  i  atom  Zn  to  Zn++ 
corresponds  to  the  change  from  2  atoms  Ag  to  2Ag+. 

A  summary  of  the  results  of  these  experiments  is  asked  for  in  Problem 
i  below. 

Solubility  of  Nitrates.  Chlorides  and  Sulfates  of  Metals.  What  conclu- 
sion can  you  draw  from  the  above  experiments  with  regard  to  the  solubility 
of  the  nitrates  of  zinc,  copper  and  silver.  State  how  you  could  prepare  the 
solid  nitrates  from  your  solutions.  Give  the  formulas.  Note  that  the  nitrates 
of  all  metals  are  soluble  in  water. 

Experiment.  To  a  small  portion  of  your  solution  containing  Ag+  add 
a  few  drops  HC1  solution;  to  another  portion  add  a  few  drops  H2SO4  solu- 
tion. Repeat  these  experiments  with  your  solution  containing  Cu++.  What 
conclusions  can  you  draw  with  regard  to  the  solubility  of  AgCl,  Ag2SO4, 
CuCL  and  CuSO4  in  water?  Note  that  nearly  all  the  chlorides  and  sulfates 
of  metals  are  soluble  in  water.  There  are  only  a  few  exceptions  in  each 
case.  Give  an  example  of  a  difficultly  soluble  sulfate. 

Preparation  of  Sulfates  from,  Nitrates  and  Chlorides.  Suggest  a  method 
of  preparing  solid  copper  sulfate  from  copper  nitrate  based  on  the  difference 
in  volatility  of  HNO,  and  H2SO4.  Try  your  method  with  your  solution 
which  contains  Cu++,  NO3~  and  H+  (which  is  usually  spoken  of  as  a  solution 
of  copper  nitrate  and  nitric  acid).  Collect  some  crystals  of  copper  sulfate  by 
evaporating  the  final  solution  to  a  very  small  volume,  and  letting  it  cool 
slowly.  Wash  the  crystals  with  a  very  little  water  and  save  them. 

Suggest  a  similar  method  of  preparing  solid  zinc  sulfate  from  a  solution 
of  zinc  chloride.  How  would  you  test  whether  the  final  product  is  free  from 
chloride  ? 

The  problem  of  preparing  a  soluble  chloride  or  nitrate  from  a  sulfate 
will  be  considered  later.  Suggest  a  method  now  if  you  can. 

Conversion  of  Soluble  Chlorides  into  Nitrates,  and  Nitrates  into  Chlor- 
ides. Experiment.  Mix  2  cc.  concentrated  HNO3  solution  and  5  cc.  concen- 
trated HC1  solution,  and  let  the  mixture  stand.  Is  there  any  evidence  of 
chemical  action?  The  same  reaction  takes  place  when  a  dilute  solution  of 
the  two  acids  is  evaporated  to  a  small  volume.  Do  not  attempt  to  write  the 
equation.  Both  acids  are  destroyed,  and  this  reaction  may  be  used  to  remove 
completely  chloride  or  nitrate  from  a  solution  by  adding  a  concentrated  solu- 
tion of  the  other  acid  in  large  excess  and  evaporating  almost  to  dryness. 

Experiment.  Dissolve  about  0.5  g.  zinc  in  hydrochloric  acid,  evaporate 
the  solution  to  a  small  volume,  add  some  concentrated  HNO3  solution,  and 
evaporate  the  mixture  almost  to  dryness  on  the  porch.  Test  the  residue  for 

30 


chloride,   and  repeat  the  treatment  with  concentrated  nitric  acid,  if  necessary. 

How  would  you  reconvert  the  zinc  nitrate  into  zinc  chloride? 

Water  of  Crystallisation.  Note  the  appearance  of  the  copper  sulfate 
crystals  you  have  prepared.  Examine  also  the  large  crystals  in  the  bottle  of 
copper  sulfate  in  the  laboratory.  The  formula  is  CuSO4  •  5H2O,  and  the  water 
present  in  the  solid  is  called  "water  of  crystallization".  Many  other  salts, 
such  as  crystalline  zinc  sulfate,  also  contain  water  of  crystallization. 

Experiment.  Heat  a  small  quantity  of  CuSO4  •  5H2O  in  a  porcelain  evap- 
orating dish,  and  note  the  change  in  its  appearance.  The  residue  is  anhy- 
drous copper  sulfate,  CuSO4.  Write  the  equation  for  the  reaction.  When 
the  dish  is  cool  add  a  few  drops  of  water.  What  reaction  will  take  place 
when  water  vapor  is  passed  over  anhydrous  CuSO4  at  a  low  temperature? 
Under  what  conditions  will  there  be  an  equilibrium? 

Give  another  example  of  an  equilibrium  involving  two  solids  and  a  gas. 
How  does  the  pressure  of  the  gas  at  equilibrium  vary  with  the  temperature? 

Problems,  i.  Summarize  what  you  know  about  the  action  of  acids  on  Na, 
K,  Ca,  Zn,  Cu,  Ag  and  any  other  metals  discussed  in  the  lectures  by  dividing 
the  metals  into  the  three  following  classes : 

1.  Those  which  react  readily  with  water.     Hydrogen  is  evolved. 

II.  Those    which    react    readily   with    HC1    or   H2SO4    solution,    but   not 
with  water.     Hydrogen  is  evolved. 

III.  Those  which  dissolve  readily  in  HNO3,  but  not  in  HC1,  H2SO4  or 

water.     H2  is  not  evolved. 
Write  the  equation  for  the  action  of  HC1  or  H2SO4  on  metallic  sodium. 

2.  How    would    you    prepare    solid    silver    sulfate    from  '  metallic    silver? 
What  weight  of  the  sulfate  could  be  obtained  from  I  gram  of  silver? 

ASSIGNMENTS  XVIII  TO  XX. 
CHEMISTRY  OF  CUPRIC  ION,   SILVER  ION,  AND  ZINC  ION. 

In  the  next  three  Assignments  we  shall  study  the  chemistry  of  the  ions 
Cu++,  Ag+,  and  Zn++, — especially  the  preparation  of  difficultly  soluble  com- 
pounds, methods  of  dissolving  them,  and  the  equilibria  involved. 

While  doing  this  work  collect  the  information  necessary  to  complete  the 
following  table,  performing  additional  experiments  when  necessary.  This 
table  is  intended  to  summarize  the  action  of  various  reagents  on  the  different 
positive  ions.  Leave  a  blank  when  there  is  no  action.  Mark  the  colors  of 
the  precipitates,  and  add  notes  in  the  table,  or  below  the  table,  to  show 
methods  of  dissolving  them.  Be  sure  that  you  can  write  the  equation  for 
each  reaction.  The  information  thus  summarized  will  be  used  in  Assignment 
XXI  in  planning  methods  of  qualitative  analysis. 
Reagents  Ca++  Cu++  Ag+  Zn++ 

Cl-  AgCl,  white 

soluble  in    ? 

OH-  Ca(OH)2, 

OH~  in  excess  moderately 

soluble. 
NH4OH 

NH4OH   in  excess 
CO3-  CaCO3 

soluble  in 

acids. 
H2S  in  acid 

solution  ? 

S™  in  alkaline 

solution  ? 

Ferrocyanide  ion 
Fe(CN).—  ? 


ASSIGNMENT  XVIII. 
HYDROXIDES  AND  OXIDES. 

Experiment.  To  solutions  containing  Cu++,  Ag+  and  Zn++,  in  separate 
test  tubes,  add  a  few  drops  NaOH  solution.  In  the  experiments  with  copper 
and  silver  continue  to  add  NaOH  solution  until  the  mixture,  after  shaking, 
reacts  strongly  alkaline  to  litmus.  Collect  the  precipitates  on  filters.  The 
precipitates  are  the  hydroxides  of  copper  and  zinc  and  the  oxide  of  silver, 
Ag2O.  Write  equations  to  show  the  formation  of  the  solids  from  the  ions. 
There  is  an  equilibrium  in  each  case.  How  will  each  equilibrium  be  affected 
by  the  addition  of  a  strong  acid? 

Test  the  action  of  nitric  acid  on  the  precipitates  just  obtained.  Com- 
pare your  equations  with  that  for  the  action  of  a  strong  acid  on  solid  calcium 
hydroxide. 

To  solutions  containing  Cu+%  Ag+,  and  Zn++,  in  separate  tubes,  add  a  few 
drops  NH4OH  solution.  Do  not  add  excess. 

Experiment.  Collect  some  zinc  hydroxide  on  a  filter.  Treat  a  portion 
with  hydrochloric  acid,  and  another  portion  with  sodium  hydroxide  solution. 
All  hydroxides  and  oxides  of  metals  react  with  acids  to  from  water  (cf.  neu- 
tralization), and  the  reaction  usually  takes  place  rapidly.  Some  of  them,  as 
in  the  case  of  zinc,  also  react  with  strong  bases  to  form  water,  and  the 
reaction  again  corresponds  to  neutralization.  Compare 

H2ZnO2  (solid)   +  2OH-  =  2H2O  +  ZnO,~ 
and  H2CO,  +  2OH~     =  2H2O  +  CO3~. 


ZnO,r~   is    zincate   ion,    and   Na2ZnO2    is   sodium   zincate.   What   reaction   takes 
place  when  Zn++  is  treated  with  excess  OH~? 

State  if  any  other  hydroxide  of  a  metal  which  reacts  with  a  strong  base 
in  a  similar  manner  has  been  discussed  in  the  lectures,  and  write  the  equation. 

Experiment.  Collect  some  Cu(OH)2  on  a  filter,  wash  it  once  with  water, 
and  heat  some  of  it  in  a  porcelain  dish.  Cupric  oxide  has  been  formed.  In 
order  to  determine  if  this  is  a  reversible  reaction,  allow  the  dish  to  cool  and 
add  water.  Also  heat  a  mixture  of  cupric  hydroxide  and  water  to  boiling. 
What  reaction  would  take  place  if  NaOH  solution  were  added  to  a  solution 
of  copper  nitrate  at  ioo°? 

The  hydroxides  of  all  metals  except  the  alkali  metals  are  decomposed  into 
the  oxide  and  water  vapor  when  the  dry  solids  are  heated  strongly.  In  most 
cases  the  oxides  will  remain  unchanged  in  contact  with  water,  ZnO  is  an 
example.  In  some  cases,  however,  the  reaction  is  reversible.  What  happens 
when  water  is  added  to  calcium  oxide?  Under  what  conditions  will  there 
be  an  equilibrium? 

The  oxides  of  the  alkali  metals,  as  Na2O,  react  violently  with  water 
Write  the  equation.  Solid  NaOH  melts  without  decomposition. 

Problems,  i.  Write  equations  for  the  action  of  HC1  solution  on  calcium 
oxide,  on  ferrous  and  ferric  oxides  (FeO  and  Fe2O3),  and  on  ferrous  and 
ferric  hydroxides. 

2.  How  would  you  prepare  solid  copper  nitrate   (a)   from  cupric  hydrox- 
ide,  (b)   from  cupric  sulfate? 

3.  Give   three   examples,   as   widely   different   as   possible,   in   which   there 
is  an  equilibrium  involving  two  solids  and  a  gas.     In  each  case  state  how  the 
pressure  of  the  gas  at  equilibrium  alters  with  the  temperature. 

32 


ASSIGNMENT   XIX. 
COMPLEX   IONS  CONTAINING  AMMONIA. 

The  ions  of  certain  metals  have  the  power  of  forming  compounds  with 
ammonia,  NH,.  The  ammonia  is  supplied  by  adding  NH4OH  solution. 
There  is  an  equilibrium  between  the  ion  of  the  metal,  NH3  or  NH4OH  and 
the  complex  ion ;  and  when  NH4OH  is  present  in  excess  the  concentration  of 
the  ion  of  the  metal  is  often  very  small.  Some  examples  will  be  studied  in 
Assignment  XIX. 

Experiment.  Collect  some  copper  hydroxide  on  a  filter,  wash  it  with  a 
little  water,  and  treat  it  with  6  normal  NH4OH  solution.  Give  reasons  why 
the  reaction  cannot  be  similar  to  that  of  OH"  on  zinc  hydroxide.  The  deep 
blue  substance  is  Cu(NH3)4++.  How  is  the  equilibrium  between  solid 
Cu(OH).:  and  its  ions  affected  by  the  addition  of  NH4OH?  Which  solution 
has  the  smaller  concentration  of  Cu++,  a  saturated  solution  of  Cu(OH)2  or 
the  solution  containing  the  complex  ion? 

From  a  consideration  of  the  equilibrium 

Cu++  +  4NH4OH  =  Cu(NH3)4++  -f  4H2O 

predict  what  will  happen  when  the  solution  is  acidified  with  nitric  acid.     Test 
your  conclusions. 

Silver  and  zinc  also  form  complex  ions  with  amonia,  Ag(NH3),+  and 
Zn(NH,)4". 

Experiment.     Prepare    some   silver   oxide   and    zinc   hydroxide,   and   test   the 
action  of  excess  NH.OH  solution. 

Predict  what  will  be  the  effect  of  treating  a  silver  chloride  precipitate 
with  NH4OH  solution.  Test  your  conclusions  by  an  experiment.  Acidify  the 
final  solution  with  nitric  acid. 

In  addition  to  the  alkali  metals  and  alkaline  earth  metals  there  are  many 
metals  which  do  not  form  complex  ions  with  amonia.  Give  examples  if 
you  can. 

Problems,  i.  Write  equations  for  the  reaction  between  Cu++  and 
NH4OH  (i)  when  a  few  drops  NH4OH  solution  are  added,  and  (2)  when 
excess  NH4OH  is  added. 

2.  Give  at  least  one  example  of  the  formation  of  a  complex  ion  of  a 
metal,  other  than  an  ammonia  complex. 

ASSIGNMENT   XX. 
CARBONATES,   SULFIDES,   FERROCYANIDES. 

Experiment.  Try  the  action  of  Na2CO3  solution  on  Cu++,  on  Ag+,  and  on 
Zn++.  In  each  case  collect  the  precipitate  on  a  filter,  wash  it  with  water,  and 
test  a  portion  for  carbonate.  Predict  the  action  of  nitric  acid  solution  and  of 
ammonium  hydroxide  solution  on  these  precipitates,  and  test  your  prediction 
by  experiments  with  HNO3  and  NH4OH  solutions. 

Predict  the  action  of  NH4OH  solution  on  calcium  carbonate,  and  test  your 
answer  by  an  experiment. 

Experiment.  To  solutions  containing  small  amounts  of  Cu++,  Ag+,  and 
Zn++,  in  separate  experiments,  add  i  cc.  6  normal  H2SO4;  dilute  each  solution 
to  about  20  cc.,  and  pass  H2S  gas*  into  it  until  the  liquid  is  saturated  with  the 
gas.  To  determine  this,  close  the  end  of  the  test  tube  or  flask,  shake 
thoroughly,  and  test  the  odor  cautiously.  In  the  experiment  with  Zn++  and 
H2S  add  5  cc.  2  normal  sodium  acetate  solution,  and  again  saturate  with  H2S 
gas.  Collect  the  precipitates  on  separate  filters. 

*  Caution.      H^S   is   poisonous.   Work   with   it  on   porch,   and  do  not   breathe   it. 

33 


Predict  the  action  of  0.3  normal  H2SO4  on  each  precipitate.  Treat  por- 
tions of  each  precipitate  with  0.3  normal  H2SO4  and  with  2  normal  H2SO4.  Is 
zinc  sulfide  more  or  less  soluble  than  copper  sulfide?  Give  your  reasoning. 

To  solutions  containing  Cu(NH,)4++,  Ag(NH3)2+,  and  Zn(NH.,)/+,  in 
separate  experiments,  add  (NH4)2S  solution.  Repeat  the  experiments,  using 
H2S  gas  instead  of  (NH4)2S  solution.  State  in  each  case  which  has  the 
smaller  concentration  of  the  ion  of  the  metal,  a  solution  containing  the  complex 
ion,  or  a  solution  saturated  with  the  sulfide.  State  also  which  is  the  less 
soluble,  Cu(OH),  or  CuS,  Ag,O  or  Ag2S,  Zn(OH)2  or  ZnS. 

Predict  the  action  of  H2S  on  mixtures  of  copper  carbonate  and  water, 
silver  carbonate  and  water,  and  zinc  carbonate  and  water.  Test  your  answers 
by  experiments.  State. in  each  case  which  is  the  less  soluble,  the  carbonate  or 
the  sulfide. 

Test  the  action  of  H2S  on  solutions  of  nitric  acid  of  different  concentra- 
tions, O.i  normal,  2.0  normal,  and  6  normal.  Saturate  each  solution  with  H2S 
gas,  and  heat  the  mixture  almost  to  boiling.  The  principal  reaction  is 

3H2S+2NO3-+2H+=3S(  solid  )+2NO+4H2O. 

Compare  this  equation  with  that  for  the  action  of  nitric  acid  on  metallic  Zn 
or   Cu. 

Collect  some  CuS  on  a  filter,  transfer  it  to  a  casserole,  add  2  normal  HNO-, 
and  boil  the  mixture.  Filter,  and  test  the  filtrate  with  excess  NH4OH.  The 
residue  collected  on  the  filter  is  sulfur;  the  dark  color  is  due  to  a  little  CuS 
enclosed  in  the  sulphur.  Complete  the  equation 

3CuS  (solid)  +2NO3-=3S  (solid)  +2NO. 
Repeat  the  experiment,  using  silver  sulfide  instead  of  copper  sulfide. 

Experiment.  To  very  dilute  solutions  containing  Cu++,  Ag+,  and  Zn++, 
respectively,  add  a  few  drops  potassium  ferrocyanide  solution,  K4Fe(CN)0. 
Repeat  these  experiments  in  the  presence  of  (i)  a  little  acetic  acid,  (2)  a  little 
hydrochloric  acid.  A  ferrocyanide  frequently  furnishes  a  characteristic  final 
test  for  an  ion,  on  account  of  the  color ;  but  ferrocyanides  are  seldom  used  in 
making  separations. 

Experiment.  Determine  which  is  the  most  delicate  test  for  copper  by 
preparing  very  dilute  solutions  of  Cu++,  e.g.  N/iooo,  N/io,ooo,  etc.,  and  testing 
separate  10  cc.  portions  with  NH4OH,  with  H2S  and  with  K4Fe  (CN)G. 

Problems,  i.  Hydrogen  sulfide  in  solution  is  a  weak  dibasic  acid  which, 
like  H2CO3,  ionizes  in  two  stages.  Review  the  neutralization  reactions  of  car- 
bonic acid  (Assignment  XIV),  and  write  equations  for  the  neutralization  of 
(i)  H2S  with  one  mol  OH~,  (2)  H2S  with  excess  OH-,  and  (3)  HS~  with 
OH".  Write  the  equation  for  the  hyd'rolysis  of  sulfide  ion,  S~~. 

2.  Is  it  correct  to  state  that  all  difficulty  soluble  salts  of  a  weak  acid  dis- 
solve when  a  solution  of  any  strong  acid  is  added?     Explain,  giving  examples. 

3.  What  is  a  "basic  salt?"     Give  an  example  of  a  basic  carbonate. 

Note.  Leave  your  two  clean  sample  bottles  at  the  office,  labelled  with  your  name 
and  desk  number,  and  marked  No.  3  and  No.  4,  in  order  that  two  unknowns  may  be  ready 
for  you  at  the  next  laboratory  period. 

ASSIGNMENT   XXI. 
REVIEW.     QUALITATIVE   ANALYSIS. 

Test  what  you  know  of  the  chemistry  of  cupric  copper,  silver  and  zinc,  as 
studied  in  the  preceding  Assignments,  by  preparing  from  memory  a  summary 
for  each  metal  showing 

I  the  formulas  of  the  ions   (including  the  complex  ions)  ; 

II  the   readily  soluble  and  moderately  soluble  compounds; 

III     the  difficulty  soluble  compounds  noting  which  is  the  least  soluble. 

34 


Check  each  item  by  reference  to  your  laboratory  notes,  and  correct  your 
mistakes. 

Complete  the  table  given  just  before  Assignment  XVIII,  and  use  this  table 
in  planning  how  to  analyze  the  various  solutions  of  salts  referred  to  below. 
Perform  additional  experiments  whenever  you  are  not  certain  that  your  methods 
are  satisfactory. 

Give  a  method  of  detecting  silver  in  a  solution  which  may  contain  Cu++, 
Ca++  and  Na+.  How  could  you  separate  silver  from  a  solution  containing  any 
of  these  ions? 

Give  two  methods  for  each  of  the  following  separations :  (a)  Cu++  f rom  Zn++, 
(b)  Cu++  from  Ca++.  (c)  Cu++  from  Na+,  (d)  Zn++  from  Ca++,  (e)  Zn++  f rom  Na+. 

A  solution  contains  Ag+,  Cu++,  Zn++,  Ca++,  and  Na+.  Devise  a  series  of 
operations  which  would  enable  you  to  prepare  a  compound  of  each  element 
from  a  single  portion  of  the  solution.  Experiment.  Test  your  method  with  a 
solution  containing  all  of  these  ions,  and  apply  additional  tests,  when  necessary, 
to  prove  that  each  of  these  ions  was  present  in  the  original  solution.  Such  a 
series  of  operations  is  a  scheme  of  qualitative  analysis  for  these  ions. 

A  knowledge  of  the  positive  ions  present  in  a  solution  often  enables  the 
conclusion  to  be  drawn  that  certain  negative  ions  cannot  be  present  in  appreci- 
able amounts.  What  negative  ions  need  not  be  tested  for  when  the  following 
are  present  at  moderate  concentrations?  (a)  H+,  (b)  Ca++,  (c)  Zn++,  (d)  Cu++, 
(e)  Ag+. 

Experiment.  Test  for  nitrate  ion.  NO3~  To  about  2  cc.  of  the  solution 
to  be  tested  add  excess  ferrous  ammonium  sulfate  (FeSO4:(NH4),JSO.1)  solu- 
tion, or  add  the  solid  salt,  filter  if  there  is  a  precipitate,  hold  the  test  tube  in  a 
slanting  position  and  pour  carefully  down  the  side  (from  a  small  beaker)  3  to 
5  cc.  concentrated  sulfuric  acid.  The  concentrated  acid  sinks  to  the  bottom  of 
the  test  tube  and  a  dark  brown  ring  forms  on  its  surface  when  nitrate  is 
present.  This  brown  substance  decomposes  when  the  solution  is  heated. 

Analysis  3  and  4.       Test  for 

H+,  Ag+,  Cu++,  Zn++,  Ca++,  Na+,  K+,  NH4+, 
OH-,  C1-,  N03-,  SO4-   CO3~,  S-. 

When  an  ion  is  found  to  be  present,  try  to  decide  whether  it  is  present  in  a 
large  amount,  in  a  small  amount,  or  as  a  mere  trace. 

In  the  second  term  A.  A.  Noyes'  "Qualitative  Analysis,"  new  edition, 
1914,  will  be  used;  it  is  for  sale  at  the  Co-operative  Store,  price  $1.50.  The 
new  edition  is  the  same  as  that  used  in  1914-15,  but  differs  greatly  from  that 
used  in  1913-14. 


35 


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